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MAIN 


B   M   57M 


LABORATORY 
OUTLINES 


IN 


PHYSICAL 
CHEMISTRY 


BY 

T.  R.   BRIGGS 


ITHACA,   NEW   YORK 
1920 


LABORATORY 
OUTLINES 

IN 

PHYSICAL 
CHEMISTRY 


BY 

T.  R.  BRIGGS 


ITHACA,  NEW  YORK 
1920 


LABORATORY  OUTLINES 

IN  PHYSICAL  CHEMISTRY  ' 


INTRODUCTION 


It  is  the  purpose  of  this  course  to  acquaint  the  student  with  some 
of  the  factors  governing  physical  and  chemical  change  and  to  enable 
him  to  recognize  these  factors  and  to  measure  their  intensity  by  their 
effects.  Painstaking  accuracy  is  not  required  in  most  of  the  experi- 
ments, which  have  been  designed  primarily  to  illustrate  principles 
and  to  encourage  intelligent  thinking.  It  is  believed  that  work  of 
this  kind  proves  more  interesting  and  stimulating  to  the  average 
student  than  do  the  more  tedious  and  exact  measurements  carried 
out  commonly  in  laboratories  of  physical  chemistry. 

Completion  of  the  work  of  this  course  entitles  the  student  to  three 
hours  of  University  credit  per  term,  two  of  which  are  given  for 
experiments  performed  satisfactorily  in  the  laboratory  and  one  for 
written  reports  based  upon  these  experiments.  The  following 
Laboratory  Outlines  describe  the  work  to  be  accomplished  during 
the  year,  though  certain  of  the  experiments  may  be  omitted  at  the 
discretion  of  the  Professor  in  charge.  No  work  of  a  similar  nature 
done  elsewhere  at  another  college  or  university  is  required  to  be 
repeated,  provided  the  work  be  submitted  to  the  Professor  in  charge 
for  his  approval. 

LABORATORY 

In  performing  the  majority  of  these  experiments,  students  are  to 
work  in  groups  of  two.  Partners  are  to  be  chosen  at  the  first  labora- 
tory period  and  this  partnership  is  to  be  maintained  throughout  the 
year  so  far  as  possible.  It  is  absolutely  essential,  however,  that  both 
partners  work  in  cooperation  on  the  same  experiment.  Independent 
work  on  different  experiments  in  a  given  group  will  not  be  permitted. 

Since  this  course  is  introductory  in  nature,  the  student  is  not  given 
the  most  delicate  instruments  or  the  purest  materials.  The  apparatus 
supplied  will  nevertheless  be  found  quite  sufficient  for  the  require- 
ments of  these  experiments.  When  the  student  has  determined  how 
closely  his  calculations  must  be  made,  he  can  readily  ascertain  the 
allowable  error,  how  carefully  his  measurements  must  be  made  and 
what  degree  of  delicacy  he  must  look  for  in  his  measuring  instruments. 
All  burettes  and  pipettes  should  be  calibrated  according  to  the 
methods  of  Experimental  Group  I  and  should  be  cleaned  in  chromic- 

1 


sulf)lno-40  nyi<l  'mxaMV  before  using.  The  use  of  dirty  or  poorly 
assembled  apparatus  will  not  be  tolerated. 

All  special  apparatus  must  be  returned  clean  and  dry  and  should 
never  be  locked  away  in  a  desk  except  by  special  permission. 

Little  attention  is  given  in  the  lectures  in  Physical  Chemistry 
(Course  50)  to  the  methods  of  experimental  physical  chemistry. 
Reference  should  therefore  be  made  constantly  to  the  Laboratory 
Manuals  and  to  other  reference  books  in  the  Chemical  Library. 
Before  commencing  work  on  any  experiment,  the  directions  should 
be  read  and  a  clear  idea  of  the  principle  involved  should  be  obtained. 

The  student  should  supply  himself  with  a  suitable  laboratory 
notebook  in  which  his  own  observations  are  to  be  neatly  recorded  at 
the  time  of  performing  the  experiment.  Recording  observations  on 
loose  sheets  of  paper  will  not  be  permitted.  Notebooks  are  to  be 
submitted  to  the  Instructor  for  approval  before  entering  upon  work 
in  this  course. 

When  making  measurements,  the  student  is  urged  to  compute  the 
results  so  far  as  possible  in  the  laboratory  at  the  time  the  work  is 
being  done  and,  if  feasible,  to  plot  rough  curves  on  cross-section 
paper.  On  the  completion  of  each  Experimental  Group,  the  labora- 
tory notes  are  to  be  submitted  to  the  Instructor  for  inspection  and 
approval  before  writing  the  final  report.  No  report  will  be  accepted 
unless  this  is  done.  . 

REPORTS 

Each  report  should  include  a  description  and  discussion  of  all  work 
completed  in  the  laboratory  together  with  answers  to  all  questions 
and  problems  appearing  in  the  Laboratory  Outlines.  Reports 
should  be  written  in  ink  and  on  one  side  of  the  paper  only,  and 
should  be  enclosed  in  a  "Department  of  Chemistry"  cover.  Care 
should  be  taken  to  describe  the  experiments  in  the  order  in  which 
they  appear  in  the  Laboratory  Outlines. 

In  writing  the  reports,  the  general  outline  given  below  should 
be  followed: 

(1)  Purpose  of  the  experiment  and  theory  illustrated. 

(2)  Apparatus  and  manipulation. 

(3)  Experimental  data  and  curves. 

(4)  Discussion. 

At  the  time  of  inspecting  the  laboratory  data,  the  Instructor  will 
assign  a  date  on  which  the  written  report  is  due.  A  deduction  of  2 
per  cent  per  diem  will  be  made  for  unexcused  lateness  in  submitting 
reports.  All  reports  are  to  be  handed  in  on  or  before  the  day  of  the 
final  examination  in  Course  50.  After  inspection  the  reports  will  be 
returned  to  the  student.  If  "double  checked"  the  report  is  accepted 
as  written.  If  "single  checked"  it  is  returned  for  correction  and 
should  be  resubmitted  with  corrections  not  later  than  one  week  after 
its  return.  When  a  report  is  received  by  the  Instructor  he  will  make 
a  note  to  that  effect  on  the  Bulletin  of  Reports  posted  in  the  labora- 
tory. Students  are  requested  to  consult  this  bulletin  and  to  notify 
the  Instructor  of  any  mistakes  or  omissions. 

A  term  grade  of  "Incomplete"  will  be  given  in  Course  51  if  at  the 
end  of  the  term  all  the  reports  have  not  been  handed  in  and  accepted. 


STANDARD  REFERENCES  IN  PHYSICAL  CHEMISTRY 
General  Texts.  Abbreviation 

Arrhenius:  Theories  of  Chemistry  (1907)  Arrhenius 

Bigelow:  Theoretical  and  Physical  Chemistry  (1912)  Bigelow 

Getman:  Outlines  of  Theoretical  Chemistry  (2d  ed.  1918)  Getman 
Hildebrand:  Principles  of  Chemistry  (1918)  Hildebrand 

Jones:  Elements  of  Physical  Chemistry  (4th  ed.)  1915  Jones 

Kremann  (Potts) :  Application  of  Physico-Chemical  Theory  (1913) 

Kremann-Potts 

Lehfeldt:  A  Textbook  of  Physical  Chemistry  (1899)  Lehfeldt 
Lewis:  A  System  of  Physical  Chemistry,  3  vols.  (1916-1918)  Lewis 
Lincoln:  Physical  Chemistry  (1918)  Lincoln 

Nernst  (Tizard):  Theoretical  Chemistry  (7th  ed.  1916)  Nernst 
Ostwald:  Lehrbuch  der  allgemeinen  Chemie  (1891-1902)  Lehrbuch 
Ostwald  (Morse):  The  Fundamental  Principles  of  Chemistry  (2d 

ed.  1917)  OFF 

Ostwald  (Walker  and  Taylor) :  Outlines  of  General  Chemistry  (2d 

ed.  1912)  OO 

Senter:  Outlines  of  Physical  Chemistry  (1911)  Senter 

van't  Hoff  (Lehfeldt):  Lectures  in  Theoretical  and  Physical 

Chemistry  (1898)  VHL 

Walker:  Introduction  to  Physical  Chemistry  (8th  ed.  1920)  Walker 
Washburn:  Principles  of  Physical  Chemistry  (1915)  Washburn 

Laboratory  Manuals. 

Biltz  (Hall,  Blanchard) :  Laboratory  Methods  of  Inorganic  Chemistry 

(1909)  Biltz 

Biltz  (Jones,  King) :     Practical  Methods  of  Determining  Molecular 

Weights  (1899)  BJK 

Ewell:     Physical  Chemistry  (1909) 

Findlay:     Practical  Physical  Chemistry  (1917)  F 

Getman:     Laboratory  Exercises  in  Physical  Chemistry  (1908)         G 
Gray:     Manual  of  Practical  Physical  Chemistry  (1.914) 
Lamb:     Laboratory  Manual  of  General  Chemistry  (1916)         Lamb 
Ostwald  (Walker) :    Physico-Chemical  Measurements  (1894)      OW 
Stabler:    Arbeitsmethoden  usw,  3  vols.  (1913)  Stabler 

Traube  (Hardin):     Physico-Chemical  Methods  (1898)  T 

Physical  and  Chemical  Tables. 

Biedermann:     Chemiker  Kalender  (annual)  Kalender 

Castell-Evans:     Physico-Chemical  Tables  (1902) 
Landolt-Bornstein-Roth :    Tabellen  (1912)  LBR 

Tables  Annuelles  de  Constantes  (1910—) 

Methods  of  Calculation — Problems. 

Knox:     Physico-Chemical  Calculations  (1916) 

Mellor :     Higher  Mathematics  for  Students  of  Chemistry  and  Physics 

(1902) 

Partington:  Higher  Mathematics  for  Chemical  Students  (1911)  — 
Prideaux:  Problems  in  Physical  Chemistry  (1912)  Prideaux 

3 


journals. 

Zeitschrift   fur   physikalische   Chemie    (1887—)    Zeit.  Phys.  Chem. 
Journal  of  Physical  Chemistry  (1896—)  Jour.  Phys.  Chem. 

Journal  de  Chimie  Physique  (1903 — )  Jour.  Chim.  Phys. 

Journal  of  the  American  Chemical  Society  (1879 — ) 

Jour.  Am.  Chem.  Soc. 
Journal  of  the  Chemical  Society  (London)  (1849 — •) 

Jour.  Chem.  Soc. 
Abstract  Journals. 

Abstract  Journal  of  the  American  Chemical  Society  (1907 — ) 
Abstracts  of  the  Journal  of  the  Chemical  Society  of  London 
Chemisches  Centralblatt  (1856—) 
Science  Abstracts  (Chemistry  and  Physics) 

PRELIMINARIES 

1.  Check  apparatus  in  desk. 

2.  Make  wash  bottle.     Use  1000  cc.  flask  in  desk. 

3.  Prepare  cleaning  mixture  as  follows: 

Dissolve  50  grams  of  powdered  commercial  Na2Cr2O?  in  about  200 
cc.  of  warm  water.  After  cooling  this  solution,  add  to  it,  slowly  and 
with  constant  stirring,  300  cc.  of  concentrated  H2SO4  (commercial). 

Keep  in  a  500  cc.  wide  mouth  bottle,  for  cleaning  grease  from  glass 
vessels. 


EXPERIMENTAL  GROUP  I 

CALIBRATION    OF    VOLUME    MEASURING    APPARATUS 

The  following  group  of  experiments  is  designed  to  give  practice  in 
testing  and  calibrating  the  volume  measuring  apparatus, supplied  to 
you  in  your  equipment.  For  accurate  work  apparatus  as  supplied 
by  the  maker  should  never  be  regarded  as  correctly  graduated  unless 
accompanied  by  the  certificate  of  the  United  States  Bureau  of  Stand- 
ards or  of  the  German  Reichsanstalt. 

Discussion. 

The  best  method  of  procedure  is  to  take  a  liquid  whose  specific 
volume  is  known  accurately  and,  completely  filling  with  it  the  appara- 
tus to  be  tested,  to  determine  the  weight  of  the  liquid  either  contained 
or  delivered.  In  most  cases  the  liquids  chosen  are  water  and 
mercury. 

Since  bodies  usually  expand  on  being  heated,  it  is  necessary  in 
calibrating  to  make  the  determinations  at  the  same  temperature 
as  that  at  which  the  apparatus  is  to  be  used.  Instead  of  doing  this, 
however,  one  may  calculate  the  volume  changes  due  to  temperature 
variations  and  may  introduce  the  necessary  corrections.  Such  cor- 
rections are  absolutely  essential  when  the  volume  of  the  apparatus 
is  large  (flasks,  etc.). 

For  accurate  work  the  calibrating  liquid  must  be  pure  and  its 
surface  free  from  contaminating  impurities  affecting  the  surface 
tension  and  hence  the  shape  of  the  meniscus.  All  volumes  are  to  be 
read  from  the  meniscus,  using  a  suitable  background  (white  or 
black). 

References. 

Read  Bulletin  U.  S.  Bureau  of  Standards,  4,  553  (1908)  or  an 
abstract  ©f  this  article  in  Mahin:  Qualitative  Analysis,  140  (1914). 
Note  carefully  units  of  capacity;  milliliter;  Mohr  units;  parallax 
and  its  avoidance;  cleaning  apparatus;  error  due  to  surface  con- 
tamination; outflow  time  and  drainage;  limit  of  error  for  burette; 
tables  for  calculation,  etc.  Cf.  also  Foulk:  Quantitative  Analysis, 
79  (1910);  OW,82;  F,  29. 

EXPERIMENT  1 
Calibration  of  Burettes 

Calibrate  a  50  cc.  burette  following  the  procedure  recommended 
by  Richards.  The  following  is  quoted  from  the  original  article  by 
Richards:  Jour.  Am.  Chem.  Soc.,  22,  149  (1900). 

"In  the  original  description  of  this  process  it  is  assumed  that  the 
calibrator  delivers  exactly  an  integral  number  of  cubic  centimeters, 


but  if  a  few  instruments  only  are  to  be  calibrated,  it  is  both  trouble- 
some and  expensive  to  secure  such  a  precise  instrument.  We  have 
found  it  convenient  to  use  a  calibrator  of  any  size,  and  in  parallel 
columns  to  compare  its  multiples  with  the  actual  readings  of  the 
burette.  The  capacity  of  this  calibrator  is  most  conveniently 
obtained  in  the  following  manner:  Suppose  that  as  a,  mean  of 
several  comparisons  it  has  been  found  that  sixteen  fillings  of  the 
calibrator  correspond  to  49.53  cc.  on  a  given  burette,  .  .  .  The 
burette  is  now  refilled  and  exactly  this  amount  of  pure  water  is  run 
into  a  weighed  flask,  with  all  the  precautions  which  would  be  used 
in  an  actual  titration.  The  weight  of  the  water  gives  by  appropriate 
calculation  the  true  volume  of  sixteen  fillings  of  the  calibrator. 
Suppose  this  was  found  to  be  49.44  cc.  ;  then  the  volume  of  the  cali- 
brator as  it  is  actually  used  in  a  calibration  must  be 


ID 

The  differences  between  the  successive  readings  of  the  burette  and 
the  successive  numbers,  3.09,  6.18,  9.27,  .  .  .  etc.,  give  at  once 
the  errors  of  the  graduation  of  the  tube  at  these  intervals.  These 
differences  or  corrections  may  be  plotted  on  a  diagram  in  which  the 
ordinates  are  volumes  and  the  abscissas  corrections.  The  correction 
to  be  applied  for  50  cc.  is  obviously  -0.09  cc." 

Notes.  Allow  the  burette  to  drain  for  two  minutes  before  making 
a  reading.  See  precaution  23  below  under  Expt.  2.  Clean  the  burette 
with  cleaning  mixture  until  the  "film  of  water  wetting  the  interior, 
will  remain  continuous  for  at  least  five  minutes"  (Bureau  of  Stand- 
ards requirement).  Results  are  of  no  value  if  grease  is  present. 

Reduce  all  weights  to  weights  in  vacuo. 

EXPERIMENT  2 
Calibration  of  Pipettes 

Calibrate  a  pipette  to  deliver  10  cc.  at  room  temperatures 
(18°-25°).  Follow  the  procedure  described  in  laboratory  manuals 
suchasF,32;  OW,84. 

Notes. 

To  secure  uniform  delivery  in  case  of  burettes  and  flasks  see  Pro- 
ceedings of  American  Chemical  Society,  21  (1904). 

"Certain  precautions  will  be  taken  to  secure  uniform  delivery. 

"18.  All  such  -apparatus  will  be  made  so  clean  internally  that  the 
film  of  water  wetting  it  will  remain  continuous  for  at  least  five 
minutes. 

"21.  Pipettes  with  one  mark  will  be  held  vertical  with  the  delivery 
orifice  touching  the  side  of  the  receiving  vessel  during  the  free  outflow 
and  for  fifteen  seconds  thereafter. 

"23.  From  burettes,  after  the  desired  volume  shall  have  been 
taken,  the  suspended  drop  will  be  removed  with  a  glass  rod  and  the 
reading  will  be  taken  at  the  end  of  two  minutes." 

6 


Note  carefully  that  the  delivery  orifice  of  a  pipette  must  be  of  such 
a  size  that  the  free  outflow  shall  last  not  more  than  two  minutes  and 
not  less  than 

12  seconds,  if  capacity  is  not  more  than  10  cc. 
15         "       "         "         lies  between  10  and  50  cc. 
20         "       "  "         "      50  and  100  cc. 

30  is  more  than  100  cc. 

EXPERIMENT  3 
Morse-Blalock  Bulb  and  Flask 

Calibrate  a  Morse-Blalock  bulb  and  flask,  the  bulb  to  deliver 
exactly  500  cc.  at  20°,  the  flask  to  hold  500  cc.  at  20°  C. 

References. 

Morse:  Exercises  in  Quantitative  Chemistry,  84  (1905) ;  Mahin: 
Quantitative  Analysis,  155  (1914) ;  also  article  in  Am.  Chem.  Jour. 
16,  479  (1894)  or  in  Olsen:  Quantitative  Analysis,  236  (1910). 

The  following  table  will  be  found  very  helpful  in  calibration  work. 
In  it  is  given  the  true  volume  of  one  apparent  gram  of  water  when  the 
latter  is  weighed  in  the  air  with  brass  weights. 

Volume  of  one  apparent  gram  of 
Temperature  (°C.)  water  (cc.) 


10  1.0014 

11  1.0015 

12  1.0016 

13  1.0017 

14  1.0018 

15  1.0019 

16  1.0021 

17  1.0023 

18  1.0024 

19  1.0026 

20  1.0028 

21  1.0030 

22  1.0033 

23  1.0035 

24  1.0037 

25  1.0040 


EXPERIMENTAL  GROUP  II 

VAPOR  DENSITY 

The  following  group  of  experiments  is  designed  to  afford  practice 
in  determining  molecular  weights  by  measuring  the  density  of 
vapors.  The  method  employed  was  introduced  by  Victor  Meyer  and 
makes  use  of  the  principle  of  air  displacement.  Before  commencing 
experimental  work,  study  the  method  carefully,  since  success  requires 
skilful  and  intelligent  manipulation. 

References. 

BJK,6-33;    F,49;   T,39;   OW,  101;  G,  30. 

Weiser:     Jour.  Phys.  Chem.,  20,  532  (1916): 

Nernst:     253  (1911)  for  measurements  at  high  temperatures. 

Turner :     Molecular  Association,  6-21  (1915) . 

Young:     Stoichiometry  (2d  ed.  1918). 

EXPERIMENT  1 
Molecular  Weight  from  Vapor  Density 

Determine  the  vapor  density  and  molecular  weight  of  an  unknown 
liquid.  Use  either  (a)  Victor  Meyer  apparatus  or  (b)  the  Weiser 
modification.  See  Instructor.  Calculate  and  report  molecular 
weight.  Check  results  before  reporting. 

Notes. 

Use  water  as  the  heating  liquid.     Boil  rapidly  and  steadily. 

It  is  a  good  plan  to  cork  the  jacket  to  insure  more  even  heating. 
The  cork,  of  course,  must  be  notched  to  permit  the  steam  to  escape. 

Use  the  earthenware  burner  guard  to  protect  the  flame  from  drafts. 
This  will  insure  steady  boiling. 

The  inner  tube  must  be  cleaned  and  dried  after  each  determination. 
Dry  by  blowing  in  air  from  the  blast,  using  a  long  delivery  tube 
reaching  to  the  bottom  of  the  inner  tube.  Pass  air  through  a 
CaCh  tube  or  tower.  The  air  in  the  apparatus  must  be  dry  at  the 
beginning  of  each  run. 

The  bottom  of  the  inner  tube  must  be  covered  with  mercury,  clean 
sand,  or  glass  wool  to  protect  it  against  breaking. 

It  is  essential  that  vaporization  should  take  place  as  rapidly  as 
possible.  If  it  takes  place  slowly,  diffusion  and  condensation  of  the 
vapor  on  the  upper  and  cooler  parts  of  the  tube  may  occur.  The 
volume  of  air  displaced  should  be  read  as  soon  as  bubbles  cease  to  pass 
over  into  the  collecting  eudiometer. 

Better  results  are  obtained  by  protecting  the  outer  jacket  from 
draughts.  Cover  the  outer  cylinder  with  asbestos  paper.  The  inner 
tube  should  not  extend  far  above  the  cork  at  the  top  of  the  heating 

8 


jacket.  The  air  displaced  by  the  vapor  of  the  liquid  must  be  at  the 
same  temperature  as  the  vapor  displacing  it.  Explain  why  this  is 
necessary. 

Do  not  attempt  to  start  this  experiment  until  you  understand  the 
operation  of  the  apparatus  and  know  the  reasons  for  the  many  and 
important  precautions. 

Take  the  following  readings  during  each  determination: 

(1)  Weight  of  sample. 

(2)  Final  volume  of  air  displaced. 

(3)  Barometer  reading  and  barometer   temperature. 

(4)  Temperature  of  water  and  temperature  of  air  surrounding 

eudiometer  tube.     These  temperatures  should  be  the  same. 

(5)  Height  of  water  column  in  eudiometer.  « 

Calculate  the  molecular  weight  of  the  unknown  substance  from  the 
above  data. 


EXPERIMENTAL  GROUP  III 

LIQUIDS  AND  LIQUID  MIXTURES 

The  purpose  of  this  group  of  experiments  is  to  study  some  of  the 
interesting  properties  of  liquids  and  liquid  mixtures,  with  special 
attention  to  volume  changes,  refractive  indices  and  viscosity. 

References. 

Dunstan  and  Thole:     The  Viscosity  of  Liquids  (1914). 
Kuenen:  Verdampfung  und  Verflussigung  (1906). 
LeBas:     Molecular    Volumes    of    Liquid    Chemical    Compounds 
(1915). 

Smiles:     Chemical  Constitution  and  Physical  Properties  (1910). 
Turner:    Molecular  Association  (1915). 
Young:     Stoichiometry  (1918). 

EXPERIMENT  1 
Change  of  Volume  and  Temperature  on  Mixing  Liquids 

Reference.     Kuenen:     Verdampfung  und  Verflussigung,  142. 

Part  1.  Mix  54  cc.  of  water  and  46  cc.  of  alcohol.  Measure 
temperature  change  and  also  change  in  volume.  Have  water  and 
alcohol  at  same  temperature  before  mixing  and  read  temperature  to 
1/5  degree  centigrade.  Obtain  thermometer  from  Instructor. 

Part  2.  Mix  equal  parts  by  volume  of  carbon  disulphide  and 
acetone.  Proceed  as  before. 

This  experiment  illustrates  the  fact  that  unexpected  and  profound 
internal  changes  often  accompany  the  mixing  of  two  liquids. 


EXPERIMENT  2 

Refractive  Index  of  Liquid  Mixtures 

The  refractive  index  of  ordinary  glass  is  1.54;  that  of  benzene,- 1 .51 ; 
while  carbon  bisulphide  has  a  refractive  index  of  1.64  for  the  same 
wave  leqgth  of  light.  One-  can  prepare  a  mixture  of  benzene  and 
carbon  bisulphide  having  the  same  refractive  index  as  glass  for  a  given 
wave  length.  The  glass  practically  disappears  as  the  refractive  index 
of  the  solution  approaches  that  of  the  glass.  Explain.  H.  G.  Wells 
has  made  fantastic  use  of  this  principle  in  his  "Invisible  Man." 

In  a  test-tube  place  about  5  cc.  of  CeHe.  Add  CS2  until  a  clean 
glass  rod,  when  dipped  into  the  mixture,  becomes  invisible. 

This  experiment  illustrates  the  fact  that  "the  properties  of  liquid 
mixtures  are  often  not  widely  different  from  the  algebraic  sum  of  the 
properties  of  the  constituents." 

10 


EXPERIMENT  3 

Relative   Viscosity   of  Benzene   and   Water 
Procedure.     F,  83;   OW,  162;   etc. 

"Having  thoroughly  cleaned  a  viscosity  tube,  introduce  into  the 
larger  bulb,  by  means-of  a  pipette,  a  known  volume  of  water,  recently 
boiled  and  allowed  to  cool,  sufficient  to  fill  the  bend  of  the  tube  and 
half,  or  rather  more  than  half,  of  the  large  bulb. 

"Fix  the  viscosity  tube  in  the  thermostat  and  after  allowing  ten  to 
fifteen  minutes  for  the  temperature  of  the  tube  and  the  water  to 
become  constant,  attach  a  piece  of  rubber  tubing  to  the  narrower 
limb  of  the  viscosity  tube  and  suck  up  the  water  to  above  the  upper 
mark.  Then  allow  the  water  to  flow  back  through  the  capillary  and 
determine  the  time  of  outflow  by  starting  the  stop  watch  as  the 
meniscus  passes  the  upper  mark.  Repeat  the  measurement  four  or 
five  times  and  take  the  mean  of  the  determinations.  If  the  time  of 
outflow  is  about  100  seconds,  the  different  readings  should  not  deviate 
from  the  mean  by  more  than  0.1  to  0.3  second.  Greater  deviations 
point  to  a  soiled  capillary  tube. 

"The  viscosity  tube  must  now  be  dried  and  an  equal  volume  of  pure 
benzene  introduced  into  the  tube  in  place  of  water.  Readings  are 
made  as  in  the  case  of  water."  F,  87. 

The  density  of  benzene  at  20°  and  water  at  20°  are  given  below.  The 
viscosity  of  benzene,  relative  to  that  of  water  at  20°  is  calculated  by 
means  of  the  formula: 

Relative  viscosity  of  benzene  =  Time  x  density  of  benzene 

Time  x  density  of  water 

Notes. 

Use  a  large  beaker  (1500  cc.)  as  a  water  bath  and  to  insure  a 
constant  temperature  keep  well  stirred.  The  compressed  air 
furnishes  an  excellent  means  of  stirring. 

It  is  important  to  keep  the  temperature  constant  because  the 
viscosity  changes  rapidly  with  the  temperature  (about  2%  per 
degree). 

Record  the  temperature  frequently.  It  is  not  sufficient  to  work  at 
room  temperature;  the  temperature  must  be  that  specified  in  the 
directions. 

See  that  the  viscosimeter  is  immersed  far  enough  to  cover  the 
upper  bulb. 

Part  1.  Determine  the  relative  viscosity  of  water  and  benzene 
at  20°. 

Part  2.     Repeat  at  40°  C. 

These  two  experiments  show  the  influence  of  temperature  on  the 
viscosity.  The  densities  follow: 

Liquid  20°  C.  40°  C. 

Water  0.9982  0.9920 

Benzene  0.8790  0.8600 

11 


EXPERIMENT  4 

Viscosity  of  Mixtures  of  Ethyl  Alcohol  and  Water 

Find  the  relative  viscosity  (water  as  standard)  of  the  following: 

Absolute  ethyl  alcohol  and  mixtures  containing  80,  60,  50,  40, 
and  20  parts  of  alcohol  in  100  parts  by  weight  of  alcohol  and  water. 

Work  at  20°  C.  ±0.1°. 

The  viscosimeter  must  be  clean.  It  is  a  good  plan  to  rinse  thor- 
oughly with  the  mixture  whose  viscosity  is  to  be  measured.  Always 
employ  the  same  volume  of  alcohol- water  mixture  in  the  viscosimeter. 

Record  the  temperature  during  each  determination. 

Draw  a  curve  with  viscosity  as  ordinates  and  composition  as 
abscissas. 

The  following  data  will  be  required: 

Parts  Alcohol  in  Mixture  Density  (20°  C.) 

0  0.9983 

20  0.9688 

40  0.9351 

50  0.9140 

60  0.8913 

80  0.8437 

100  0.7895 

This  experiment  is  another  illustration  of  the  fact  that  in 
many  cases  the  properties  of  a  mixture  are  unexpectedly  different 
from  the  properties  of  the  pure  constituents.  Compare  with  Experi- 
ment 1  above.  Could  one  use  viscosity  measurements  as  a  means  of 
determining  the  alcohol  content  of  alcohol-water  mixtures? 


EXPERIMENT  5 
Relative  Viscosity  of  Unknown 

Determine  the  relative  viscosity  of  an  unknown  solution,  using 
water  as  standard.  The  Instructor  will  supply  the  unknown  solution 
and  will  state  the  temperature  at  which  to  work,  and  the  density  of 
the  solution. 

EXPERIMENT  6 
Specific  Gravity  Flotation 

If  a  mixture  of  dry  sawdust  and  iron  filings  is  thrown  into  water, 
the  sawdust  will  float  and  the  iron  filings  will  sink,  the  two  being 
separated  by  means  of  a  liquid  whose  specific  gravity  lies  between 
those  of  the  mixed  solids. 

Employing  this  principle,  separate  the  mixture  of  two  solids 
which  is  found  on  the  shelf.  Use  as  the  liquid  the  solution  formed 
when  HgI2  dissolves  in  an  excess  of  KI. 

Specific  gravity  of  the  lighter  solid  is  2.05. 

Specific  gravity  of  the  heavier  solid  is  3.60. 

12 


The  specific  gravity  of  the  liquid  should  be  about  midway  between 
these  values. 

The  solution  is  made  by  adding  saturated  KI  solution  to  the 
saturated  HgCl2  solution  on  the  shelf.  Avoid  large  excess  of  KI. 
Put  in  a  test  tube  and  shake  violently,  for  two  or  three  minutes.  See 
T,  17;  also  Stahler  1,  626  (1913).  Danger;  Mercuric  chloride  is 
extremely  poisonous ! 

Compare  flotation  of  this  type  with  froth  flotation  now  used  on  so 
large  a  scale  for  the  concentration  of  sulphide  ores.  See  Mineral 
Industry,  24,  807  (1915);  Megraw:  The  Flotation  Process  (1917). 


13 


EXPERIMENTAL  GROUP  IV 

VAPOR  PRESSURE 

The  following  group  of  experiments  is  designed  for  the  purpose  of 
studying  and  measuring  the  pressure  exerted  by  the  vapor  phase 
when  in  equilibrium  with  a  pure  liquid  or  with  a  liquid  mixture. 

References. 

Kuenen:     Verdampfung  und  Verflussigung  117,  129  (1906). 
Young:     Stoichiometry  260  (1908). 

Also  see  references  under  "Distillation,"  Experimental  Group  VII; 
Papers  by  Smith  and  Menzies  in  Jour.  Am.  Chem.  Soc.  (1910 — ). 

Apparatus. 

One  heavy-walled  test  tube  150  mm.  long,  25  mm.  ext.  diameter, 
fitted  with  two-hole  rubber  stopper;  Chapman  water  suction  pump 
(large  size) ;  Mercury  manometer ;  Y-tube ;  two  glass  stopcocks ; 
pressure  tubing,  etc.  Instead  of  test  tube  and  rubber  stopper  a 
special  glass  stoppered  test  tube  may  be  employed  with  better  results. 

Procedure. 

Refer  to  the  diagram  of  apparatus.  The  heavy  test  tube  A,  the 
vaporization  vessel  containing  the  liquid  under  investigation,  is 
immersed  in  a  constant  temperature  water  bath.  The  two  stopcocks 
are  placed  at  Pr  and  P2.  B  serves  as  a  trap.  The  remainder  of  the 
diagram  requires  no  explanation. 

First  assemble  the  whole  apparatus,  connect  with  manometer  and 
pump  and,  closing  Pt  and  opening  P2,  test  the  apparatus  for  leaks. 
If  the  pump  is  working  properly  a  "vacuum"  of  2-3  cm.  should  be 
obtained.  Read  the  barometer  in  the  balance  room  and  from  this 
reading  subtract  the  reading  obtained  on  your  manometer.  The 
pressure  in  A  should  not  exceed  35  mm.  and  should  remain  constant 
on  closing  P2. 

Place  in  A  the  liquid  whose  vapor  pressure  is  to  be  measured. 
Use  10-20  cc.  Replace  stopper  completely,  submerge  the  whole 
test  tube  in  the  constant  temperature  bath  and  proceed  with  the 
measurement.  Close  Px  and  open  P2.  Gently  agitate  the  liquid  in 
A  by  shaking  the  test  tube  back  and  forth;  this  will  tend  to  prevent 
bumping  during  vaporization.  When  the  liquid  begins  to  vaporize 
or  to  boil  slightly,  close  P2  and  continuing  the  shaking  to  hasten 
equilibrium,  read  the  manometer  when  the  latter  remains  constant. 
Again  open  P2  and  vaporize  for  an  instant.  For  the  second  time 
close  P2  and  read  the  manometer  as  before.  When  repeated  vaporiza- 
tions of  very  short  duration  fail  to  cause  an  appreciable  change  in  the 
manometer  readings,  and  the  difference  in  the  mercury  levels  in  the 
two  arms  of  the  manometer  reaches  a  maximum  and  is  constant, 
subtract  this  difference  from  the  height  of  the  barometer.  The  value 
so  calculated  is  the  vapor  pressure  of  the  liquid  in  A. 

14 


Notes. 

Be  sure  that  the  suction  pump  is  clean  and  is  operating  properly. 
Be  on  guard  against  violent  bumping  when  the  liquid  in  A  boils. 
Shake  A  to  prevent  this  and  to  hasten  adjustment  of  thermal  equili- 
brium between  the  liquid  in  A  and  the  water  bath.  (Experiment: 
place  some  ether  in  A  and  connect  with  the  vacuum  pump.  Note 
temperature  of  ether). 

Make  certain  that  all  air  has  been  removed  from  A  before  taking 
the  final  reading  of  the  manometer.  The  vaporizing  process  should 
remove  the  air. 

In  dealing  with  solutions  vaporize  no  more  than  is  just  necessary 
to  remove  air.  Boiling  or  vaporizing  a  solution  almost  always 
changes  the  composition  of  both  liquid  and  vapor.  Explain. 

Compare  the  vapor  pressures  so  determined  with  those  given  in  the 
tables  in  LBR.  The  results  with  pure  liquids  should  not  be  more 
than  2  per  cent  in  error. 

EXPERIMENT  1 
Vapor  Pressure  and  Composition 

Part  A.  Ascertain  the  vapor  pressure- com  position  relations  in 
the  system  ethyl  alcohol  and  benzene,  a  pair  of  consolute  liquids. 
Temperature  20°  C. 

Measure  the  vapor  pressure  of  the  following:  Pure  ethyl  alcohol; 
benzene;  mixtures  of  alcohol  and  benzene  containing  10,  25,  32,  50, 
75,  and  90  parts  of  alcohol  by  weight  in  100  parts  of  mixture.  Den- 
sity benzene  =  0.88;  alcohol  (absolute)  =  0.78  at  20°  C.  Mix,  using 
burettes. 

Plot  a  curve  as  you  proceed  with  the  determinations ;  pressures  as 
ordinates,  compositions  as  abscissae. 

Note  that  the  results  with  the  solutions  are  approximate  only, 
because  the  vaporization  process,  especially  if  prolonged,  causes  the 
composition  of  the  liquid  phase  to  change  and  thereby  to  be  different 
from  the  composition  of  the  original  mixture.  For  accurate  work  the 
composition  of  the  liquid  at  the  end  of  the  experiment  should  be 
determined.  The  method  as  described  approaches  sufficiently  close 
to  the  more  accurate  method  to  enable  the  student  to  obtain  the 
characteristic  pressure-composition  diagram. 

Part  B.  Ascertain  the  vapor  pressure-composition  relations  in 
the  system  acetone  and  chloroform.  Temperature  20°  C. 

Measure  the  vapor  pressure  of  the  following:  Acetone;  chloro- 
form; mixtures  of  acetone  and  chloroform  containing  15,  25,  40,  50, 
60,  75  and  85  parts  of  acetone  in  100  parts  of  mixture.  Density 
acetone  =  0.80;  chloroform  =  1.52. 

Plot  a  curve  as  you  proceed.     Compare  with  Part  A. 

Note.    Do  either  Part  A  or  Part  B  as  assigned. 

EXPERIMENT  2 
Two  Liquid  Layers 

Determine  the  vapor  pressure  of  pure  ethyl  acetate  at  20°  C. 
Look  up  the  vapor  pressure  of  water. 

15 


Determine  the  vapor  pressure  of  the  following  mixtures  of  ethyl 
acetate  and  water  containing  25,  50  and  75  parts  of  ethyl  acetate  in  100. 
Explain  your  results.     Density  ethyl  acetate  =  0.923. 

EXPERIMENT  3 

Raoult's  Law 

Part  A.  Determine  the  lowering  of  vapor  pressure  when  5  g.  of 
naphthalene  are  dissolved  in  20  g.  of  acetone.  From  this  calculate 
the  molecular  weight  of  naphthalene,  using  Raoult's  formula: 


a) 


p         '    N+n 

p  =  vapor  pressure  of  pure  acetone  ;  p'  =  vapor  pressure  of  solution  ; 
n  =  gram  molecules  of  naphthalene  ;  N  =  gram  molecules  of  acetone. 

Part  B.  Determine  the  molecular  weight  of  nitrobenzene  (8g.) 
in  ether  (25  cc.).  Density  ether  =  0.73;  nitrobenzene  =  1.2. 
Determine  vapor  pressure  of  ether  separately. 

Note.     Do  either  (A)  or  (B)  as  assigned. 

Part  C.     Optional  Experiment;    Menzie's  Method. 

Using  Menzie's  apparatus  (see  Instructor)  determine  the  molecular 
weights  of  naphthalene  or  nitrobenzene.  Reference:  Bigelow,  320. 

Notes. 

Look  up  the  vapor  pressures  of  naphthalene  and  nitrobenzene  at 
20°  C.  Are  these  solutes  volatile  at  this  temperature? 

How  is  the  lowering  of  vapor  pressure  made  use  of  in  Burger's 
method  of  determining  molecular  weights  when  very  small  amounts 
of  substances  are  available?  Jour.  Chem.  Soc.,  85,  286  (1904); 
Chamot:  Chem.  Microscopy,  216  (1915). 

EXPERIMENT  4 
Vapor  Pressure  of  Aqueous  Solutions 

For  this  work  connect  the  apparatus  with  the  rotary  vacuum 
pump,  protecting  the  latter  from  moisture  by  means  of  a  tower  or 
tube  containing  anhydrous  calcium  chloride.  Temperature  20°  C. 

Determine  the  vapor  pressure  of  (a)  water  (b)  5  per  cent  cane 
sugar  solution  (c)  30  per  cent  cane  sugar  solution  (d)  a  solution 
containing  enough  calcium  chloride  to  be  equimolecular  with  the 
sugar  solution  in  (b).  Explain  all  results. 

EXPERIMENT  5 
Vapor  Pressure  (Dissociation  Pressure)  of  Salt  Hydrates 

A  crystalline  salt  hydrate  will  effloresce  (dissociate)  when  exposed 
to  the  air  if  the  partial  pressure  of  water  vapor  in  the  air  is  less  than 
the  dissociation  pressure  of  the  hydrate.  Measure  the  dissociation 
pressure  of  Glauber's  salt  (sodium  sulphate  decahydrate),  correspond- 
ing to  the  reaction: 

Na2S04.  10H20  =  Na2S04  +  10H2O 

Reference.     Findlay:     The  Phase  Rule,  83  (1904). 

Compare  with  Experiments  1  and  2  of  Experimental  Group  VIII. 

16 


EXPERIMENTAL  GROUP  V 

ELEVATION  OF  THE  BOILING  POINT 

This  group  of  experiments  deals  particularly  with  the  changes 
produced  in  the  boiling  point  when  a  soluble,  non-volatile  substance 
is  'added  to  a  pure  solvent.  The  differences  between  electro- 
lytes and  non-electrolytes  are  emphasized  and  explained^  by  means 
of  the  theory  of  electrolytic  dissociation.  Molecular  weights  are 
determined  by  the  so-called  "boiling  point"  method. 

References.  BJK,  141-197;  F,138;  OW,184;  G,77;  T,97.  Read 
Bigelow,  317. 

EXPERIMENT  I 

Preliminary  Experiment  to  Illustrate  the  Difference  Between  an 
Electrolyte  and  a  Non-Electrolyte 

Typical  non-electrolytes — sugar,  urea,  etc. 

Typical  electrolytes — NaCl,  KC1,  etc. 

This  experiment  is  to  be  performed  by  groups  of  students  working 
together  under  the  direct  supervision  of  the  Instructor. 

Place  two  clean  graphite  electrodes  in  350  cc.  of  distilled  water 
contained  in  a  500  cc.  beaker.  Connect  electrodes  to  110-volt 
alternating- current  circuit  in  series  with  a  lamp-bank  resistance. 

Short  circuit  the  current  across  the  electrolyzing  cell  and  observe 
the  brightness  of  the  lamps.  The  brightness  is  roughly  a  measure 
of  the  current  flowing.  Then  pass  the  current  through  the  distilled 
water  and  observe  again  the  brightness  of  the  lamps. 

Substitute  350  cc.  of  tap  water  for  distilled  water.  What  do  you 
observe? 

Finally  test  in  order  5  g.  of  the  following  substances  dissolved  in 
350  cc.  of  distilled  water:  Sodium  chloride,  mercuric  chloride, 
cane  sugar,  and  acetic  acid.  Carefully  wash  the  graphite  electrodes 
after  each  solution  has  been  tested. 

What  can  you  say  regarding  the  power  o'f  the  above  solutions  to 
conduct  the  electric  current?  Is  a  good  electrolyte  always  an 
inorganic  salt,  and  are  all  inorganic  salts  good  .electrolytes? 

Note.  A  very  neat  experiment  similar  to  the  above  is  described 
in  Lamb,  33. 

EXPERIMENT  2 
Elevation  of  the  Boiling  Point 

Procedure.  Cf.  Bigelow,  319,  Walker  and  Lumsden:  Jour. 
Chem.  Soc.,  73,  502  (1898). 

For  this  experiment  a  modified  and  simple  form  of  the  Lands- 
berger  apparatus  for  vapor  heating  is  employed.  The  import- 
ant features  are  three,  viz,  vapor  (steam)  generator,  boiling  chamber 

17 


(fitted  with  Beckmann  thermometer,  outer  jacket  and  exit  tube) 
and  suitable  condenser.  See  the  diagram. 

The  steam  generator  should  be  operated  at  constant  speed  and 
without  "bumping".  To  ensure  this,  protect  the  burner  with  an 
earthenware  guard  and  add  pumice  generously  to  the  water  in  the 
round-bottom  flask.  Do  not  change  the  rate  of  boiling  during  a 
given  run  and  do  not  shut  off  or  move  the  burner  under  any  circum- 
stance. 

Set  a  Beckmann  thermometer  for  the  boiling  point  of  water. 
Make  sure  that  the  mercury  is  low  on  the  scale.  Handle  with  care 
the  delicate  and  expensive  thermometer. 

Start  the  generator  boiling  and,  when  ready,  connect  to  the 
boiling  chamber  containing  the  solvent.  The  boiling  chamber  should 
be  well  insulated  thermally.  This  may  be  done  by  using  a  Dewar 
tube  (thermos  vacuum  bottle),  by  surrounding  the  tube  with  the 
vapor  of  the  solvent  as  in  the  McCoy  apparatus  (which  see),  or  by 
slipping  the  large  test  tube  serving  as  boiling  chamber  into  a  wide- 
mouth  bottle,  fitting  snugly,  and  closing  the  annular  space  at  the 
neck  with  felt  or  cotton  wool.  The  delivery  tube  for  the  steam 
should  reach  to  the  bottom  of  the  boiling  chamber  and  the  Beckmann 
thermometer  should  be  immersed  far  enough  to  submerge  the  bulb. 
Weigh  the  dry  test  tube  so  that  the  weight  of  the  solution  whose  boil- 
ing point  is  measured  may  be  determined. 

Place  pure  water  in  the  boiling  chamber  and  boil  with  steam. 
When  the  mercury  reaches  a  steady  position  on  the  scale,  take  a  series 
of  ten  consecutive  readings  at  intervals  of  ten  seconds.  Use  a  read- 
ing glass  (to  be  obtained  from  Instructor).  The  readings  should  not 
fluctuate  by  more  than  one  of  the  smallest  divisions  (0.01  °C.). 
Read  the  barometer  before  and  after. 

Then,  without  interrupting  the  boiling,  disconnect  the  steam  line 
from  the  boiling  chamber,  lift  the  cork  holding  the  thermometer, 
and  drop  into  the  water  in  the  tube  a  weighed  quantity  of  solute. 
Determine  the  new  boiling  temperature. 

Stop  the  run,  remove  the  large  test-tube,  cool  and  weigh.  Cal- 
culate the  weight  of  water  in  the  solution.  Using  the  formula 

M=520^  (1) 

compute  the  molecular  weight  of  the  dissolved  solute. 
Beckmann  Differential  Thermometers. 

Cf.  F,  129-133. 

"Some  thermometers  have  scales  which  allow  the  adjustment  of 
the  zero  point  when  desired.  One  kind  has  a  scale  which  may  be 
screwed  up  or  down  from  the  top.  Another  kind  permits  a  change  in 
the  volume  of  mercury.  The  Beckmann  is  of  the  latter  type.  This 
thermometer  has  at  the  upper  end  of  the  capillary  a  mercury  reservoir 
which  allows  one  to  decrease  or  increase  the  actual  amount  of  mer- 
cury in  the  bulb  and  capillary  thread.  To  decrease  the  mercury  in 
the  bulb,  the  bulb  is  heated  until  the  needed  amount  of  mercury 
appears  in  the  reservoir  as  a  globule,  then  a  sharp  tap  with  the  hand 
will  separate  it,  if  the  thermometer  is  held  in  an  upright  position.  It 

18 


is  apparent  then  that  the  temperature  of  the  bath  should  be  higher 
than  the  required  zero  reading  by  the  number  of  degrees  correspond- 
ing to  the  length  of  thread  which  is  not  required." 

A  good  Beckmann  thermometer  should  fulfill  these  requirements: 

(1)  The  upper  and  lower  mercury  reservoirs  should  branch  into 
the  capillary  in  a  conical  fashion. 

(2)^  Mercury  should  be  clean. 

(3)  Thermometer  should  not  be  unnecessarily  clumsy. 

Part  1.     Non-Electrolytes. 

Determine  the  molecular  weights  of  urea  and  cane  sugar.  Use 
1  /20th  g.  molecule  of  each  substance.  From  your  own  data  calculate 
the  elevation  you  would  have  observed  if  the  solutions  had  contained 
exactly  500  grams  of  water.  .  How  do  these  elevations  compare 
with  each  other? 

Part  2.    Electrolytes. 

Proceed  as  in  Part  1  with  KC1  and  K2SO4.     Compare  with  Part  1. 

Part  3.     Mercuric  Chloride. 

Proceed  as  in  Part  1  with  mercuric  chloride  (poison).  Compare 
with  urea  and  sugar  and  with  KC1  and  K2SO4. 

Part  4.     Unknown. 

Determine  the  molecular  weight  of  an  unknown  substance. 
After  obtaining  good  checks  report  your  results  to  the  Instructor. 

Part  5.     Ethyl  Alcohol  as  Solvent. 

Place  absolute  alcohol  in  the  outer  compartment  and  about  6  g.  of 
absolute  alcohol  in  the  inner  compartment  of  a  McCoy  vapor  heater. 
Guard  against  fire  by  connecting  a  long  rubber  tube  to  the  side  arm. 
When  the  alcohol  has  boiled  for  some  time  close  this  rubber  tube  with 
a  pinchcock  and  heat  the  alcohol  in  the  inner  compartment  with 
alcohol  vapor.  The  inner  compartment  is  fitted  with  a  stopper  con- 
taining an  exit  tube  connected  with  a  condenser  and  a  Beckmann,  the 
bulb  of  which  is  immersed  in  the  alcohol. 

Determine  the  boiling  point  of  pure  absolute  alcohol. 

Then  add  —  molecular  weight  of  urea  to  the  alcohol  in  the  inner 

compartment  and  heat  the  solution  with  the  vapor.  When  the 
boiling  point  has  reached  a  maximum,  pour  the  contents  of  the  inner 
tube  into  a  bottle  and  determine  the  weight  of  the  solution. 

For  this  experiment  the  Beckmann  must  be  set  for  the  boiling 
point  of  absolute  alochol. 

Boiling  constant  for  ethyl  alcohol,  1170.  Calculate  the  molecular 
weight. 

Notes. 

Redetermine  the  boiling  point  of  the  pure  solvent  before  each  run. 
If  this  is  not  done  and  only  one  determination  is  made,  the  barometric 
(atmospheric)  pressure  may  change  enough  to  give  very  misleading 
results.  At  about  100°  C.  a  change  in  pressure  of  only  1  mm.  of 
mercury  produces  a  temperature  difference  in  the  boiling  point  of 
nearly  four  hundredths  of  a  degree.  Bigelow,  317. 

19 


For  a  critical  discussion  of  the  method  and  a  very  elegant  apparatus 
for  determining  the  elevation  of  the  boiling  point  read  Cottrell; 
Jour.  Am.  Chem.  Soc.,  41,  721  (1919)  and  Washburn  and  Read: 
Ibid.,  41,737  (1919). 

EXPERIMENT  3 

Lowering  of  the  Boiling-Point 
Discussion. 

A  non- volatile  solute  added  to  a  pure  liquid  always  raises  the  boil- 
ing point.  When  however  a  non- volatile  solute  is  added  to  a  mixed 
solvent  containing  two  volatile  liquids  a  depression  of  the  boiling 
point  may  be  produced  instead  of  an  elevation. 

Let  A  and  B  be  two  volatile  substances  forming  a  single  homogene- 
ous solution.  Call  A  the  solvent  and  B  the  solute.  As  B  is  added  to 
A  the  concentration  of  the  solution  increases  and  the  partial  pressure 
of  A  in  the  vapor  becomes  smaller  (Raoult's  law).  At  the  same  time 
the  partial  pressure  of  B  increases  in  the  vapor  (Henry's  law).  When 
A  is  saturated  with  B  the  solution  is  in  equilibrium  with  pure  B  and 
the  partial  pressure  of  B  in  the  vapor  is  practically  equal  to  the 
vapor  pressure  of  pure  B. 

It  follows  from  this  that,  for  a  given  concentration  of  B  in  A,  the 
greater  the  solubility  of  B,  the  smaller  is  the  partial  pressure  of  B  in 
the  vapor.  Anything  which  decreases  the  solubility  will  tend  to 
increase  the  partial  pressure  of  the  solute  in  the  vapor. 

The  solubility  of  B  in  A  may  be  made  less  by  the  addition  of  a  suit- 
able third  substance.  If  the  latter  is  non-volatile  and  soluble  both 
in  A  and  B,  it  can  affect  the  total  vapor  pressure  of  the  solution  in 
two  ways,  as  follows: 

(1)  By  decreasing  the  solubility  of  B  in  A  (or  A  in  B). 

(2)  By  dissolving  in  A  and  B. 

Influence  (1)  points  in  the  direction  of  increased  vapor  pressure 
and  may  in  fact  be  greater  than  influence  (2)  which  tends  toward  a 
lower  vapor  pressure.  (Why?)  The  total  vapor  pressure,  which 
is  equal  to  the  sum  of  the  partial  pressures  of  A  and  B,  may  thereby 
be  increased  and  the  boiling  point  depressed.  "  The  experiments 
which  follow  illustrate  the  point. 

Procedure. 

Determine  the  boiling  point  of  a  mixture  of  50  parts  alcohol  and 
50  parts  water.  Use  a  flask  and  reflux  condenser,  determining  to 
tenths  of  one  degree  with  a  special-thermometer  (not  the  Beckmann). 
Then  add  sodium  carbonate  to  the  alcohol-water  mixture  and  rede- 
termine  the  boiling  point.  Do  two  layers  appear  as  carbonate  is 
added? 

Repeat,  using  cane  sugar  instead  of  Na2CO3. 

It  may  be  found  advisable  to  add  the  Na2CO3,  or  sugar  in  several 
portions,  determining  the  boiling  point  each  time. 


20 


EXPERIMENTAL  GROUP  VI 

DEPRESSION  OF  THE  FREEZING  POINT 

The  object  of  the  following  group  of  experiments  is  the  study  and 
use  of  the  freezing  point  method  of  determining  molecular  weights. 

References.     F,  125-138;   T,  81-90;   OW,  180-184;   etc. 
Procedure. 

Apparatus:  Freezing  point  apparatus  (see  diagram) ;  Beckmann 
thermometer;  reading  glass;  etc. 

Set  the  Beckmann  for  the  freezing  point  of  water. 

As  solvent  use  10-15  cc.  of  distilled  water,  i.  e.  enough  to  cover  the 
bulb  of  the  thermometer. 

In  the  battery  jar  place  a  freezing  mixture  of  salt  and  ice.  The  ice 
must  be  pounded  fine  and  be  well  mixed  with  salt.  The  best 
temperature  for  the  freezing  bath  is  about — 5°  C.  A  lower  tempera- 
ture than  this  is  undesirable.  Record  temperature  of  freezing  mix- 
ture. See  Findlay  on  "convergence  temperature." 

In  the  freezing  mixture,  place  the  outer  tube  or  jacket,  and  in  the 
jacket,  the  inner  tube,  which  must  not  come  in  contact  with  the  walls 
of  the  outer.  The  jacket  should  be  closed  by  a  cork  through  which 
the  outer  tube  passes. 

Determine  first  the  freezing  point  of  the  solvent,  noting  the  degree 
of  undercooling  (supercooling)  and  tapping  the  thermometer  fre- 
quently to  prevent  stiction.  The  water  must  be  stirred  constantly 
to  prevent  excessive  undercooling.  Take  the  tube  out  of  the  jacket 
and  warm  in  the  hand  until  the  ice  melts.  Redetermine  the  freezing 
point.  Undercooling  should  not  exceed  1°  C. 

The  preliminary  cooling  may  be  hastened  by  placing  the  inner  tube 
directly  in  the  freezing  mixture.  Take  care  that  no  salt  from  the 
freezing  mixture  is  introduced  into  the  solution  and  dry  the  tube  very 
carefully  before  replacing  in  the  outer  tube. 

The  inner  tube  should  be  closed  by  a  cork  through  which  the  ther- 
mometer and  stirrer  pass. 

EXPERIMENT  1 

Water  as  Solvent 

Determine  the  molecular  weight  of  an  unknown  salt.  Use  about 
15  (weighed  to  O.lg.)  of  water  and  not  more  than  0.3g.  of  the  un- 
known. When  your  results  check  satisfactorily,  report  them  to  the 
Instructor.  Constant  for  water,  1860. 

21 


EXPERIMENT  2 
Benzene  as  Solvent 

Part  1 .  Determine  the  molecular  weight  of  naphthalene  or  anthra- 
cene in  benzene.  Use  about  1  /1000th  gram-molecule  of  solute  in  10  g. 
of  benzene  (thiophene  free).  Set  the  Beckmann  for  benzene  (5.5°  C.) 
and  use  ice  alone  (no  salt)  as  the  freezing  agent.  Constant  for 
benzene,  5000. 

Part  2.  Proceed  as  above  with  benzoic  acid  in  benzene.  How  do 
you  account  for  the  high  value  of  the  molecular  weight? 

For  accurate  results  the  benzene  should  be  anhydrous  and  should 
be  protected  from  moisture  in  the  air. 

EXPERIMENT  3 

To  show  how  the  Freezing  Point  of  a  Metal  may  be  affected  by  other 

Metals 

The  melting  point  is  used  as  a  criterion  of  purity,  especially  in 
organic  chemistry.  This  experiment  shows  how  one  substance  affects 
the  melting  point  of  another. 

The  fusible  alloy  is  a  mixture  of  Bi,  Cd,  Pb,  Sn. 

Place  about  0.1  g.  in  a  small  glass  tube  which  has  been  closed  at  one 
end  by  drawing  down  and  fusing.  Find  the  melting  point  of  the 
alloy  in  a  water  bath. 

Look  up  the  melting  points  of  the  metals  composing  the  alloy  in 
LBR,  190. 


EXPERIMENTAL  GROUP  VII 

DISTILLATION  OF  LIQUID  MIXTURES 

The  following  experiments  are  designed  to  illustrate  the  distillation 
of  mixtures  both  constituents  of  which  are  volatile  at  the  boiling 
point.  Particular  emphasis  is  laid  on  the  relations  existing  between 
boiling  temperature  and  the  composition  of  residue  and  distillate. 

References. 

Kuenen:     Verdampfung  und  Verfliissigung  (1906). 
Ostwald:     Fundamental  Principles  of  Chemistry  123-148. 
Rosanoff:     Jour.  Am.  Chem.  Soc.  (1909-). 
Young:     Fractional  Distillation  (1903). 
Young:     Stoichiometry  (1918). 

EXPERIMENT  1 

Hydrochloric  Acid  and  Water 
Solutions. 

(a)  1  liter  of  10  per  cent  HC1  —  i.  e.  (10  g,  HC1,  90  g.  water). 

(b)  500  cc.  30  per  cent  HC1. 

(c)  2  liters  2NNaOH,  standardized  against  N/l  HC1. 
(Use  rubber  stopper  for  reagent  bottle). 

Part  1.     Distillation  of  the  10  per  cent  Mixture 

Discussion. 

When  two  miscible  liquids  are  distilled,  the  composition  of  residue 
and  distillate  (vapor)  will  generally  differ  at  any  given  temperature 
of  ebullition  and  the  latter  will  rise  as  the  distillation  is  continued. 
The  distillate  (vapor)  will  always  be  richer  in  respect  to  the  more  vola- 
tile constituent  or,  if  the  pair  of  liquids  gives  a  mixture  of  minimum 
boiling  point  (water  and  ethyl  alcohol),  the  distillate  will  be  richer 
than  the  residue  in  respect  to  this  mixture.  If,  however,  the  pair  of 
liquids  gives  a  mixture  with  a  maximum  boiling  point  (HC1  and  water 
HNO3  and  water;  H2SO4.and  water)  the  distillate  will  be  richer  than 
the  residue  in  respect  to  either  one  of  the  pure  constitutents,  depend- 
ing upon  conditions.  What  these  conditions  are  will  be  shown  by  the 
following  experiments. 

Procedure. 

Place  500  cc.  of  the  10  per  cent  solution  in  a  liter  distilling  flask 
connected  with  condenser  and  receiver.  Place  the  thermometer  in 
vapor  and  use  ebullition  tubes  or  pumice  to  prevent  bumping. 

Before  starting  to  distill  determine  the  KC1  content  of  the  solution 
by  titration  with  standard  NaOH.  Withdraw  a  5  cc.  sample  from 
the  flask  with  a  pipette. 

23 


Distill  and  collect  the  distillate  in  a  measuring  cylinder.  Wherr 
about  30  cc.  of  distillate  have  been  collected,  remove  the  measuring 
cylinder  and  empty  it  of  its  contents  as  completely  as  possible.  Then 
collect  between  5  and  8  cc.  of  fresh  distillate,  noting  the  average 
temperature  at  which  it  comes  over.  Stop  the  distillation. 

Withdraw  a  5  cc.  sample  of  distillate  and  determine  its  HC1  con- 
tent. Next  withdraw  rather  more  than  5  cc.  of  hot  residue  in  a  flask, 
cool  and  titrate  a  5  cc.  sample. 

Again  distill;  collect  another  30  cc.;  throw  this  away  as  before 
and  collect  a  second  sample  of  5  to  8  cc.,  observing  the  tempera- 
ture. Continue  until  nearly  all  of  the  acid  has  been  distilled  over. 

Arrange  data  as  follows: 


(a) 

(b) 

(c) 

(d) 

(e) 

(f) 

Number 

Nature  of 

Volume  of 

NaOH 

Grams 

Tempera- 

of Sample 

Sample 

of  Sample 

(cc.) 

HC1  per 

ture 

(cc.) 

100  cc. 

(average) 

No.  1 

Original 

5 

6.60 

9.45 

99.5° 

No.  2 

Distillate 

5 

0.15 

0.22 

110.5° 

Residue 

5 

6.80 

9.65 

100.5° 

etc. 

etc. 

etc. 

etc. 

etc. 

etc. 

Part  2.     Distillation  of  the  30  per  cent  Mixture  (Hood). 

The  procedure  requires  modification,  since  at  the  start  nearly  pure 
gaseous  HC1  is  given  off.  Do  not  determine  the  composition  of  the 
distillate  until  the  distillation  is  nearly  finished.  Instead,  analyze 
samples  of  the  residue  at  appropriate  intervals  and  observe  the 
temperature  immediately  prior  to  withdrawing  the  samples. 

Connect  the  condenser  to  an  absorption  train  for  the  removal  of 
HC1  fumes  (adapter  dipping  below  the  water  in  beaker)  and  work 
in  hoods. 

When  the  temperature  has  reached  a  nearly  constant  value  remove 
the  absorption  apparatus  and  proceed  exactly  as  in  the  previous  case, 
analyzing  both  distillate  and  residue. 

Part  3.     Distillation  of  10  per  cent  Mixture  with  Vigreux  Column. 

Start  with  500  cc.  of  acid  mixture  in  a  round  bottom  flask  to  which 
a  long  Vigreu  column  has  been  fitted.  Place  a  thermometer  at  the 
head  of  the  column  in  the  usual  fashion,  also  a  thermometer  in  the 
vapor  in  the  flask.  Take  simultaneous  reading  of  both  thermometers 
throughout. 

Proceed  with  the  10  per  cent  solution  just  as  in  Part  1,  analyzing 
both  distillate  and  residue.  Continue  the  distillation  until  residue 
and  distillate  have  the  same  composition. 

Computations  and  Curves. 

Calculate  results  in  terms  of  grams  of  HC1  in  100  cc.  of  solution. 
On  a  single  sheet,  draw  curves  between  temperatures  as  ordinates  and 

24 


composition  as  abscissae.  Do  all  the  curves  approach  a  common 
point?  What  is  the  effect  of  the  Vigreux  column?  Explain. 

Compute  the  percentage  of  HC1  by  weight  in  the  mixture  of  maxi- 
mum boiling  point.  Consider  the  specific  gravity  of  the  mixture  to 
be  1 .1 .  Use  the  data  as  determined  by  the  experimental  curves. 

From  the  data  derive  a  formula  for  the  constant  boiling  mixture, 
assuming  that  it  is  a  definite  hydrate  of  hydrochloric  acid.  How  was 
the  simple  hydrate  theory  disproved? 

EXPERIMENT  2 
McCoy  Apparatus  and  Vapor  Heating 

Part  1.  In  the  outer  compartment  of  a  McCoy  apparatus  place 
ethyl  alcohol.  Connect  a  condenser  to  one  of  the  side  arms;  to  the 
other  a  short  piece  of  rubber  tubing  fitted  with  a  pinch  cock.  Keep- 
ing side  arm  open,  heat  the  alcohol  to  boiling.  In  the  inner  compart- 
ment place  5  or  10  cc.  of  benzene  and  close  the  tube  with  a  cork  carry- 
ing a  thermometer  dipping  into  the  benzene.  When  the  alcohol  is 
boiling  very  gently  and  evenly,  close  the  pinchcock  and  pass  alcohol 
into  the  benzene.  Read  time  and  temperature  at  intervals  of  15 
seconds.  Draw  a  curve  with  times  as  abscissae  and  temperatures  as 
ordinates. 

Precaution.  Do  not  begin  heating  with  vapor  until  the  ther- 
mometer in  the  benzene  registers  higher  than  75°  C;  then  pass  in 
alcohol  vapor  as  slowly  as  possible.  The  rate  of  heating  should  be 
kept  constant  throughout. 

Part  2.  Repeat  with  acetone  in  the  outer  compartment  and 
chloroform  in  the  inner. 

Part  3.  Repeat  with  water  in  the  outer  compartment  and  methyl 
alcohol  in  the  inner. 

Part  4.  Repeat  with  ethyl  acetate  in  the  inner  compartment  and 
water  in  the  outer.  Observe  carefully  the  formation  of  two  layers. 
Why  does  the  temperature  remain  constant  and  how  does  it  compare 
with  the  boiling  temperature  of  pure  ethyl  acetate  and  pure  water? 
Explain. 

Part  5.  Repeat  with  water  in  the  inner  compartment  and  ethyl 
acetate  in  the  outer. 

See  Experiment  4  below. 

% 

EXPERIMENT  3 
Steam  Distillation 

Take  two  1000  cc.  distilling  flasks.  In  one  place  distilled  water, 
beads  to  prevent  bumping,  and  a  thermometer  reading  to  110° 
immersed  in  the  liquid.  In  the  other  place  a  concentrated  solution 
of  NaCl  and  add  NaCl  in  large  excess.  In  this  flask  place  a  thermo- 
meter reading  to  at  least  125°  and  immerse  in  the  liquid.  See  sketch. 

25 


Boil  the  water  in  the  first  flask  and  when  the  water  is  boiling 
gently,  connect  to  the  other  flask  and  pass  steam  into  the  salt  solution . 
Note  the  temperature  in  each  flask,  making  frequent  readings. 

When  the  temperature  in  the  flask  containing  the  solution  has 
reached  a  maximum,  take  the  temperature  of  the  vapor  in  each  flask. 
Thoroughly  wash  the  thermometer  with  water  after  withdrawing 
from  the  solution,  and  again  take  the  temperature  of  the  vapor. 
Explain  the  results. 

Regarding  the  differences  observed  when  the  thermometer  is 
immersed  in  the  vapor  and  not  in  the  liquid,  see  Hite:  Am.  Chem 
Jour.  17,  510  (1895);  Sakurai:  Jour.  Chem.  Soc.,  61,  495  (1892). 


EXPERIMENT  4 

To  Map  out  the  Boiling  Point — Composition  Diagram  for  a  Binary 
Liquid  Mixture 

Determine  the  temperature  at  which  the  liquid  mixture  boils 
steadily.  Use  a  small  round-bottomed  flask  and  not  more  than  30 
grams  of  liquid  in  each  case.  The  neck  of  the  flask  should  be  fairly 
wide  and  should  be  fitted  with  a  cork  carrying  a  thermometer  and 
connected  with  a  reflux  condenser.  Place  the  thermometer  in  the 
liquid  mixtures  (chloroform-acetone  or  benzene-alcohol)  that  you 
studied  in  Experimental  Group  IV,  Experiment  1  A  or  1  B.  Having 
determined  the  boiling  point,  plot  the  values  against  the  composition. 
Compare  with  the  pressure-composition  diagram. 


2t> 


*    EXPERIMENTAL  GROUP  VIII 

DISSOCIATION 

The  following  experiments  are  designed  to  illustrate  qualitatively 
the  dissociation  of  ehemical  compounds,  either  as  the  result  of  an 
increase  in  temperature  or  as  the  result  of  dissolving  the  substance  in 
a  solvent.  Dissociation  of  the  first  type  is  called  thermal;  dissocia- 
tion of  the  second  type  is  called  electrolytic  when  ions  are  formed. 
We  have  already  studied  some  of  the  phenomena  due  to  electrolytic 
dissociation,  especially  in  Experimental  Groups  IV  and  V.  Other 
instances  of  electrolytic  dissociation  and  its  effects  will  be  studied  in 
the  Experimental  Groups  which  follow. 

References.     Solutions  and  Electrolytic  Dissociation. 

Abegg:  Die  Theorie  der  elektrolytischen  Dissociation,  Ahren's 
Sammlung  8  (1903). 

Arrhenius:     Theories  of  Solution  (1912). 
Findlay:     Osmotic  Pressure  (2nd  Ed.  1919). 
Jacques:     Complex  Ions  (1914). 
Jones :     The  Nature  of  Solution  (1917) . 
Ostwald  (Muir) :     Solutions  (1891). 
Rothmund:     Die  Loslichkeit  (1907). 

Scxidder :  Electrical  Conductivity  and  lonization  Constants  (1914) . 
Seidell:  Solubilities  of  Inorganic  and  Organic  Compounds  (1919). 
Stieglitz:  Qualitative  Analysis,  Vol.  I  (1917). 

EXPERIMENT  1 
Thermal  Dissociation  of  Nitrogen  Tetroxide 

In  a  test  tube  heat  a  small  quantity  of  Pb(NO3)2  and  pass  the 
resulting  gas  through  a  delivery  tube  into  a  test  tube  which  is 
surrounded  by  a  freezing  mixture  of  ice  and  salt. 

The  NO2  will  condense,  under  these  conditions,  as  a  bluish  green 
liquid,  N2O4.  On  removing  from  the  cooling  bath  the  colorless  gas 
N2O4  will  be  formed  first  and  on  further  heating  this  will  dissociate 
into  NO2.  Note  color  changes. 

References.  Nernst,  453  (1911).  Ostwald:  Principles  of  Inor- 
ganic Chemistry,  329  (1908). 

EXPERIMENT  2 
Thermal  Dissociation  of  Limestone 

Heat  some  powdered  marble  in  a  hard  glass  tube.  Show  that  dis- 
sociation takes  place. 

References.  Bigelow4  etc.  For  study  of  the  reaction  used  in 
lime  burning  read  Kremann- Potts:  107;  LBR,  398. 

27 


Define  "dissociation  pressure"  and  draw  a  curve  showing  how 
dissociation  pressure  changes  with  the  temperature  for  the  following 
reaction:  2NaHCO3  =  Na2CO3  +  H2O  +  CO2. 

Reference.     LBR,  398. 

* 

EXPERIMENT  3 
Electrolytic  Dissociation  and  Color 

Part  1.  Compare  the  colors  of  concentrated 'solutions  of  the  fol- 
lowing salts:  CuSO4,  CuCl2,  CuBr2.  Dilute  until  they  have  the 
same  blue  color.  Start  with  about  one  cc.  of  solution.  Explain. 

Part  2.  Add  concentrated  hydrochloric  acid  to  a  greenish-blue 
solution  of  CuCl  2.  Note  color  change.  Also  heat  some  of  the  same 
solution.  Explain. 

References.  Ostwald:  Prin.  Inorg.  Chem.,  642  (1908);  also 
Mellor:  Inorganic  Chemistry,  468  (1914). 

Part  3.     Color  changes  with  CoCl2  solutions. 
Dissolve  a  little  cobalt  chloride  in  absolute  alcohol. 
Add  two  or  three  drops  of  water  to  the  solution. 
Add  ether  to  the  solution. 
Add  water  again. 

Also  to  the  pink  solution  in  water  add  (1)  solid  magnesium  chloride ; 
(2)  concentrated  hydrochloric  acid. 

References.  Donnan  and  Bassett:  Jour.  Chem.  Soc.,  102,  939 
(1902).  Ostwald:  Prin.  Inorg.  Chem.,  623  (1908);  Nernst,  389 
(1911). 

EXPERIMENT  4 
Reactions  depending  upon  Degree  of  Dissociation 

Part  1.  Pass  chlorine  gas  into  AgNO3  solution.  Does  a  precipi- 
tate form  at  once? 

Part  2.     Add  carbon  tetrachloride  to  AgNO3  solution.     Explain. 

Part  3.  Add  chloroform  to  AgNO3  solution.  Does  a  precipitate 
form  at  first? 

Let  -the  mixture  stand  in  the  light  until  the  next  period.  Does  a 
precipitate  form  on  standing?  Explain. 


EXPERIMENT  5 
Complex  Ions 

Part  1.  To  AgNO3  solution  add  KCN  in  excess.  Test  for  silver 
with  NaCl.  Do  not  use  HCU  Beware  of  HCN  and  remember  that 
KCN  is  poisonous.  Use  hoods. 

Part  2.  To  AgNO3  solution  add  sodium  thiosulphate'in  excess. 
Test  for  silver.  Explain  results.  Cf.  Walker,  343. 

28 


EXPERIMENT  6 
Relative  Stability  of  Complex  Ions 

Part  1.     Add  KCN  in  excess  to  dilute  CdSO4.     Test  for  cadmium 
with  H2S. 

Part  2.     Add  KCN  in  excess  to  dilute  CuSO4.     Test  for  copper 
with  H2S. 

Explain.     Cf.  Walker,  343. 

Caution.     Cyanogen  is  formed  in  Part  2.     Use  hoods. 


EXPERIMENT  7 
Hydrolysis 

Part  1.  Test  KCN  and  Na2CO3  solutions  with  litmus  paper. 
Explain. 

Part  2.  Test  CuSO4  and  A12  (SO4)3  solutions  with  litmus  paper. 
Explain. 

Part  3.  Precipitate  PbSO4  completely  from  lead  acetate  solution 
by  adding  A12  (SO4)3.  Then  add  water  and  boil.  Filter  and  test 
the  filtrate  for  lead  and  aluminum. 

Precaution.  It  is  essential  to  use  very  little  Al2(SO4)3  in  excess. 
At  any  rate,  add  plenty  of  water  and  boil  thoroughly  for  several 
minutes.  Explain. 

EXPERIMENT  8 

Conductivity  and  Electrolytic  Dissociation 
Discussion. 

The  conductivity  is  the  reciprocal  of  the  resistance.  From  the 
resistance  of  a  solution,  its  conductivity  may  be  calculated.  In  this 
experiment  the  relative  resistance  of  N/10  HC1  and  N/10  CH3COOH 
is  measured  by  reading  the  current  and  voltage  across  graphite  elec- 

-p 
trodes  which  dip  into  the  solution.     From  Ohm's  law,  I   =  — '  the 

R 
resistance  may  be  computed. 

By  maintaining  the  temperature  constant,  keeping  the  electrodes 
the  same  distance  apart,  and  having  them  immersed  to  the  same 
extent,  a  rough  approximation  of  the  conductivity  of  these  two 
equivalent  acid  solutions  may  be  obtained. 

The  conductivity  of  a  solution  depends,  among  other  things,  upon 
its  dissociation.  If  two  solutions  are  of  equivalent  concentration  and 
at  the  same  temperature  and  if  both  are  placed  in  the  same  vessel  for 
measuring  the  conductivity,  the  better  conducting  solution  is  either 
more  completely  ionized  or  else  contains  the  more  mobile  (the  more 
rapidly  moving)  ions.  If  the  difference  in  conductivity  is  very  great, 
as  in  the  present  case,  the  poorly  conducting  solution  is  almost  certainly 
the  less  strongly  dissociated.  Since  both  solutions  have  the  hydrogen 
ion  in  common  and  since  the  chlorine  and  acetate  ions  are  about 

29 


equally  mobile,  the  relative  conductivity  is  here  a  very  nearly  exact 
measure  of  the  relative  ionization. 

Procedure. 

Measure  the  relative  resistance  of  N/10  HC1  and  N/10  CH,COOH 
solutions. 

Follow  the  procedure  used  in  the  experiment  which  showed  the 
distinction  between  an  electrolyte  and  a  non-electrolyte  Use 
alternating  current  and  a-c  meters. 

Look  up  the  per  cent  ionization  of  N/10  HC1  and  N/10  CH3COOH. 
Are  your  conductivity  values  proportional? 


30 


EXPERIMENTAL  GROUP  IX 

SOLUTION  AND  SOLUBILITY 

The  experiments  of  the  following  group  are  designed  to  illustrate 
the  process  of  solution,  the  properties  of  saturated  solutions,  the  cor- 
rosion or  solution  of  metals  and  the  determination  of  solubility. 

References.     See  under  Group  VIII — Dissociation. 


EXPERIMENT  1 
Quantitative  Determination  of  Solubility 

References.     F,  302;   OW,  176;   G,  234,  etc. 

The  solubility  of  a  salt  in  water  depends  chiefly  upon  the  nature  of 
the  salt  and  the  temperature.  The  rate  at  which  the  salt  dissolves 
depends  upon  the  same  factors  plus  several  others  besides,  such  as 
size  of  particles,  rate  of  stirring,  presence  of  catalysts,  and  so  forth. 

Solubility  may  be  determined  directly,  provided  the  salt  is  not  too 
slightly  soluble,  by  saturating  a  solution  with  an  excess  of  salt  at  a 
desired  temperature,  and  analyzing  a  definite  weight  or  volume  of  the 
solution. 

Determine  the  solubility  of  an  assigned  salt  at  25°  C.  Place  in  a 
bottle  an  excess  of  finely  powdered  salt,  add  water  and  shake  in  a 
thermostat  until  equilibrium  is  reached,  or  until  there  is  no  change  of 
density  between  successive  tests,  when  measured  with  a  delicate 
hydrometer.  In  a  second  bottle  place  finely  divided  salt  and  add, 
not  water,  but  a  solution  of  the  salt  saturated  at  some  temperature 
(usually  a  higher  one)  at  which  the  salt  is  more  soluble  than  it  is  at 
25°  C.  Shake  as  before  and  determine  the  density  of  the  saturated 
solution.  The  final  densities  should  be  the  same  in  both  bottles. 

Withdraw  samples  for  analysis  using  a  dry  pipette  and  a  small 
filtering  tube  to  prevent  the  entry  of  solids.  Determine  the  concen- 
tration of  the  saturated  solution  either  by  chemical  analysis,  or  by 
evaporating  a  weighed  sample  to  dry  ness  in  an  oven  or  desiccator. 
Check  results.  Determine  the  density  of  the  solution  at  25°  C.  and 
calculate  the  solubility  of  the  salt  in  grams  per  100  grams  of  solution; 
also  in  terms  of  the  "molar  fraction"  of  the  solute. 


EXPERIMENT  2 
Cryolite  and  Water 

Add  a  little  finely  powdered  cryolite  to  water  in  a  test  tube.     Does 
it  dissolve?     Explain. 

31 


EXPERIMENT  3 
Solution  and  Catalysis 

Chromic  chloride  appears  in  two  forms,  as  the  hexahydrate 
(CrCl3.  6H2O)  green  in  color,  and  as  the  anhydrous  salt  (CrCl3) 
which  is  violet.  The  anhydrous  form  appears  to  be  nearly  insoluble 
in  water  while  the  hydrate  dissolves  readily.  According  to  Moissan 
the  violet  form  dissolves  slowly  at  high  temperatures  to  a  green 
solution,  and  Ostwald  believes  that  the  apparent  insolubility  at 
ordinary  temperatures  is  due  to  the  extreme  slowness  with  which 
solution  occurs;  in  other  words,  that  the  violet  form  is  not  really  in 
equilibrium  with  water.  Drucker  under  Ostwald's  direction  showed 
that  the  violet  modification  dissolves  readily  in  the  presence  of 
chromous  chloride  (CrCl2)  in  solution,  the  latter  acting  as  a  catalyst. 

With  these  facts  in  mind  perform  the  following  tests: 

(1)  Try  to  dissolve  violet  CrCl3  in  water. 

(2)  Dissolve  some  green  hexahydrate  in  water. 

(3)  To  a  small  quantity  of  the  violet  salt  add  water  plus  a  crystal 
of  the  green  hexahydrate.     Add  a  bit  of  zinc  and  acidify  with  HC1. 
See  whether  the  violet  salt  dissolves  in  time.     Explain. 

(4)  To  th'e  violet  salt  add  a  bit  of  metallic  chromium,  then  add 
dilute  HC1.     Does  the  salt  dissolve? 

(5)  Prepare  chromous  chloride  by  dissolving  metallic  chromium 
in  dilute  HC1.     Add  this  solution  to  a  few  particles  of  the  violet  salt. 
Do  they  dissolve? 

(6)  Repeat  (d)  adding  zinc  instead  of  chromium.     Explain. 

References.  Ostwald:  Prin.  Inorg.  Chem.  615;  Mellor:  Inorg. 
Chem.  258. 

Drucker:     Zeit.  phys.  Chem.,  36,  173  (1901). 

EXPERIMENT  4 
Relative  Solubility 

Part  1.  Precipitate  PbSC>4,  let  it  settle,  wash  once  or  twice  by 
decantation,  then  add  KI  solution  to  the  residue.  Note  the  color 
change.  Then  warm  it.  What  color  change  occurs? 

Part  2.  Precipitate  AgCl,  repeat  procedure  in  (a)  using  KBr 
solution.  What  change  occurs  in  the  precipitate?  Explain. 

Part  3.     Repeat  (2)  using  KI  solution.     Walker,  356  (1913). 

Part  4.  Precipitate  AgBr,  add  KC1  solution.  Is  there  any  visible 
change?  Explain. 

Part  5.  Prove  by  simple  experiments,  which  is  the  more  soluble, 
CaSO4  or  CaCO3. 

EXPERIMENT  5 
Compound  Solvents 

Part  1.  Add  about  20  cc.  of  impure  commercial  sulphuric  acid 
to  an  equal  volume  of  water.  What  is  the  precipitate?  Explain. 

32 


Part  2.  Add  about  5  cc.  of  95  per  cent  ethyl  alcohol  to  (1)  a 
saturated  solution  of  Na2SO4  (2)  a  saturated  solution  of  Na2CO3. 
cf.  Group  V,  Experiment  3. 

Part  3.  Determine  by  experiment  qualitatively  the  effect  of 
sodium  chloride  on  the  solubility  of  phenol  in  water.  Repeat  with 
sodium  acetate  instead  of  sodium  chloride. 

EXPERIMENT  6 
Solubility  Product 
Discussion. 

When  a  salt,  dissociating  into  univalent  cations  and  anions, 
is  in  equilibrium  with  its  saturated  solution,  the  Law  of  Mass 
Action  leads  to  the  conclusion  that  the  product  of  the  concen- 
trations of  cation  and  anion  is  a  constant  for  a  given  temperature, 
provided  the  nature  of  the  solvent  undergoes  no  change.  The 
product  of  the  ion  concentrations  when  the  solution  is  saturated  is 
called  the  solubility  product.  Thus : 

[cation]  [anion]  =  Ks,  (1) 

where  the  symbols  "[cation]"  etc.,  represent  the  concentrations. 

Reference.  Stieglitz,  I  141  (read  page  142  for  criticism  of 
theory) ;  Washburn,  298. 

When  the  salt  dissociates  into  polyvalent  ions  or  into  ions  of  mixed 
valence,  the  relation  is  more  complex.  Cf.  Washburn,  301. 

It  is  possible  to  distinguish  between  two  cases,  as  follows : 

(1)  When  to  a  solution  saturated  with  a  given  solid  electrolyte 
there  is  added  a  soluble  salt  containing  a  common  ion,  the  product  of 
the  concentrations  of  cation  and  anion  momentarily  becomes  greater 
than  the  solubility  product.     The  solution  is  no  longer  in  equilibrium 
with  the  saturating  solid  salt  and  the  latter  is  precipitated,  until  new 
conditions  of  equilibrium  are  established.     These  new  conditions 
correspond  to  diminished  solubility. 

(2)  When  the  concentration  of  one  or  both  of  the  ions  produced  by 
the  saturating  solid  is  decreased  by  any  kind  of  physical  or  chemical 
reaction,  the  product  of  the  concentrations  of  cation  and  anion 
momentarily    becomes    less    than     the     solubility    product.     The 
solution  is  no  longer  in  equilibrium  with  the  solid  and  fresh  solid 
dissolves  until  new  conditions  of  equilibrium  are  established,   the 
latter  corresponding  to  increased  solubility. 

Procedure. 

Part  1.  To  a  BaCl2  solution  in  a  test  tube  add  concentrated  HC1, 
then  add  water.  Explain. 

Part  2.     Repeat,  using  a  CaCl2  solution. 

Part  3.     To  a  saturated  solution  of  NaCl,  add  concentrated  HC1. 

Part  4.  To  a  saturated  solution  of  HgCl2  add  a  saturated  solution 
of  NaCl. 

Account  for  what  happened  in  (1)  and  (3),  by  applying  the  theory 
of  the  solubility  product.  Explain  the  very  different  results  of  (2) 
and  (4). 

33 


How  might  all  these  experiments  be  explained  in  the  light  of 
Experiment  4? 

Reference.     Ostwald:     Prin.  Inorg.  Chem.,  675  (1908). 

Part  5.  Treat  some  freshly  precipitated  and  washed  AgCl  with 
(a)  Na2S2O3  solution;  (b)  with  KCN  solution  (poison);  (c)  with 
NH4OH. 

Part  6.  Treat  some  freshly  precipitated  calcium  oxalate  with 
HC1. 

Part  7.  Shake  a  little  HgO  with  a  solution  of  KI.  Note  any  color 
change.  Filter  and  test  the  nitrate  with  red  litmus. 

Part  8.  Prepare  some  Cd(OH)2  and  wash  thoroughly  with  water. 
Shake  with  water  and  test  the  supernatant  liquid  with  red  litmus. 
The  solution  should  be  neutral.  To  one-half  of  the  Cd(OH)2  add  a 
small  amount  of  KNO3  and  shake  again.  Test  the  supernatant 
liquid  with  red  litmus.  To  the  second  half  of  the  Cd  (OH)2  add  a 
little  KI,  shake  and  test  the  supernatant  liquid*  with  red  litmus. 
Explain. 

Reference.     Ostwald:     Prin.  Inorg.  Chem.,  637  (1908). 


EXPERIMENT  7 
Solubility  of  Glass 

Part  1.  Phenolphthalein  Test.  Boil  some  clean,  finely-powdered 
glass  with  water  in  a  beaker,  then  add  a  drop  of  phenolphthalein. 
Explain. 

Part  2.  Eosin  Test.  "If  a  glass  surface  is  brought  into  contact 
with  watery  ether,  it  draws  water  from  the  solution  and  gives  up 
alkali  to  it.  On  the  other  hand,  the  orange-yellow  solution  of  iod- 
eosin  in  ether  is  changed  by  the  alkali  into  red.  Mylius,  who  had 
previously  used  the  color  reaction  for  another  purpose,  has  applied  it 
to  the  practical  testing  of  glasses.  Commercial  ether  is  shaken  up 
with  water  at  ordinary  temperature  until  it  is  saturated  with  water. 
It  is  then  poured  from  the  rest  of  the  water  and  eosin  is  added  in  the 
proportion  of  0.1  g.  to  100  cc.  of  the  liquid.  The  solution  is  filtered 

"Glass  vessels  are  tested  by  pouring  in  the  solution.  The  first 
step  is  to  clean  the  surface  from  any  products  of  weathering  which 
may  adhere  to  it,  by  carefully  rinsing  with  water,  with  alcohol,  and 
lastly  with  ether.  Immediately  after  the  cleaning  with  ether,  the 
eosin  solution  is  poured  in,  the  vessel  is  carefully  closed  and  the  solu- 
tion is  allowed  some  twenty-four  hours  to  do  its  work.  It  is  then 
emptied  out  and  the  glass  rinsed  with  pure  ether.  The  surface  of  the 
glass  is  now  seen  to  be  colored  red;  and  the  strength  of  the  color 
furnishes  an  indication  of  the  susceptibility  of  the  glass  to  attack  by 
cold  water." 

Reference.  Hovestadt  (Everett) :  Jena  Glass  and  its  Scientific 
and  Industrial  Applications. 

34 


Following  these  directions  make  up  100  cc.  of  eosin  solution. 
Then  test  the  surface  of  a  new  50  cc.  beaker  and  a  new  test  tube  as 
described  above. 

Place  a  small  sample  of  powdered  glass  in  an  8-dram  vial  and  add 
some  eosin  solution. 

Note  the  color  the  powdered  glass  assumes  on  standing  twenty-four 
hours.  Note  also  the  color  of  the  walls  of  the  vial. 

If  the  powdered  glass  becomes  colored,  filter  it  and  wash  thoroughly 
with  water.  Does  the  water  remove  the  color?  Pour  off  the  water 
and  add  alcohol.  Does  the  alcohol  remove  the  color? 

Eosin  as  Indicator.  Take  a  few  cubic  centimeters  of  the  eosin 
solution  and  add  a  few  drops  of  dilute  NaOH. 

Part  3.  Tetrachlorgallein  Test.  Add  to  a  beaker  of  boiling  dis- 
tilled water  a  few  drops  of  alcoholic  tetrachlorgallein.  Continue  the 
boiling  and  observe  the  color  change.  Make  a  blank  test  with  fresh 
distilled  water. 

EXPERIMENT  8 

Corrosion  of  Metals 
Discussion. 

Many  metals  dissolve  more  or  less  readily  in  aqueous  solutions, 
appearing  in  the  solution  in  the  form  of  cations  for  at  least  a  limited 
time  and  displacing  during  this  process  an  equivalent  weight  of  some 
other  cation,  usually  hydrogen,  from  the  solution.  Thus  zinc  and 
sulphuric  acid  give  zinc  sulphate  and  hydrogen;  zinc  and  copper 
sulphate  give  zinc  sulphate  and  metallic  copper,  the  salts  and  acids 
being  in  solution.  Under  these  circumstances  the  zinc  is  said  to 
corrode. 

It  is  generally  believed  that  the  process  of  corrosion  is  electro- 
chemical in  nature.  For-example,  when  zinc  corrodes,  two  so-called 
"electrochemical"  reactions  take  place  as  follows: 

(1)  Metallic  zinc  gives  zinc  ions  plus  negative  charges,  the  latter 
being  retained  by  the  metal. 

(2)  Hydrogen  ions  in  solution  give  hydrogen  gas  plus  positive 
charges,  the  latter  neutralizing  the  negative  charges  on  the  metal. 

Represented  by  symbols,  these  reactions  may  be  written: 
+26  (1) 

+20  (2) 

If  one  adds  reactions  (1)  and  (2),  the  total  reaction  becomes 

Zn  +  2H+— ^Zn  +  +  +  H2  (3) 

It  is  interesting  to  note  that  the  anions  appear  to  play  no  part 
whatsoever. 

Applying  the  Law  of  Mass  Action  to  the  two  reactions  given  above, 
it  is  possible  to  draw  the  following  conclusions  regarding  the  rate  of 
corrosion : 

(1)  A  metal  tends  to  corrode  more  readily  in  an  aqueous  solution 
the  greater  its  "electrolytic  solution  pressure,"  i.  e.,  the  greater  the 
driving  force  of  reaction  (1)  or  the  greater  the  ion-forming  tendency 
of  the  metal. 

35 


(2)  The  smaller  the  concentration  of  the  dissolving  metal  as  ion 
in  the  solution,  the  faster  is  the  corrosion.     The'  ion  concentration 
may  be  kept  low  by  the  formation  of  complex  ions,  by  hydrolysis,  etc. 

(3)  The  greater  the  hydrogen  ion  concentration  in  the  solution  the 
faster  the  corrosion.     Other  things  being  equal,  metals  tend  to  cor- 
rode more  readily  in  acids  than  they  do  in  alkaline  solutions. 

(4)  Anything  that  reacts  with  and  removes  the  discharged  hydro- 
gen tends  to  aid  corrosion.     Oxidizing  agents  may  do  this,  in  which 
case  they  are  called  "hydrogen  depolarizers."     Note  the  part  played 
by  air  in  the  experiments;   also  the  formation  of  nitrites  in  Part  2b. 

(5)  The  absence  of  stable,  difficulty  soluble  protecting  films  (pas- 
sivity) favors  corrosion. 

(6)  Miscellaneous.     Metal    should    have    irregularities,    etc.,    in 
surface  to  aid  in  setting  up  local  "galvanic"  couples.     Also  the 
"overvoltage"  for  hydrogen  should  be  low.     These  points  belong 
properly  under  electrochemistry  and  cannot  be  discussed  here. 

All  the  conditions  favoring  corrosion  do  not  have  to  be  fulfilled 
simultaneously.  Copper  for  example  corrodes  in  aqueous  ammonium 
hydroxide  in  the  presence  of  air.  The  electrolytic  solution  pressure 
of  copper  is  very  small  and  the  hydrogen  ion- concentration  in  ammo- 
nium hydroxide  solution  is  very  slight,  but  these  conditions  which 
tend  to  prevent  corrosion  are  more  than  offset  by  the  fact  that  the 
copper  ion  concentration  in  the  solution  is  practically  zero  (complex 
Cu(NH3)2  cations)  and  air  oxidizes  the  discharged  hydrogen  under 
the  conditions  of  the  experiment.  The  reaction  as  a  whole  may  be 
written : 

Cu  +  2NH4OH  +  O— >Cu  (NH3)2  (OH)2  +  H2O 

Iron  corrodes  readily  in  moist  air.  Moisture  is  essential  inasmuch 
as  it  furnishes  the  hydrogen  ions  which  are  displaced  by  the  iron,  the 
latter  entering  the  solution  in  the  form  of  ferrous  ions.  These  are 
almost  immediately  oxidized  by  air  to  ferric  ions  which  combine  with 
the  hydroxyl  ions  of  the  water  to  form  hydrous  ferric  oxide.  The 
iron  thus  passes  from  solution  and  corrosion  is  thereby  accelerated. 
Carbon  dioxide  stimulates  corrosion  by  dissolving  in  the  film  of  mois- 
ture and  thus  increasing  the  hydrogen  ion  concentration  by  the  forma- 
tion of  H2CO3.  Air  increases  corrosion  by  removing  the  dissolved 
iron  as  explained  above  and  by  serving  as  the  hydrogen  depolarizer. 

Procedure. 

Part  1.  Solubility  of  Metals  in  Acids  and  Alkalies,  (a)  Place 
a  small  strip  of  copper  foil  in  aqueous  NH^OH  in  a  test  tube.  Shake 
thoroughly  from  time  to  time.  Note  the  color  change  and  explain. 

(b)  Experiments  with  concentrated  H2SO4. 

In  a  few  cc.  of  concentrated  H2SO4  test  the  solubility  of  cast  iron, 
iron  wire,  nickel  wire,  and  copper  wire.  Set  aside  for  an  hour. 
Dilute  the  acid  five  fold  with  water  and  repeat,  using  the  same  test 
pieces.  Dilute  the  acid  until  the  rate  of  solution  is  rapid.  Caution. 
Dilute  the  acid  properly. 

36 


(c)  Experiments  with  concentrated  HNO3. 

In  a  few  cubic  centimeters  of  concentrated  HNO3,  test  the  solu- 
bility of  iron  wire  and  nickel  wire.  Set  aside  for  an  hour.  Repeat 
with  acid  diluted  twice.  Why  are  metals  often  more  readily  attacked 
by  HNO3  than  they  are  by  HC1? 

Test  the  solubility  of  aluminum  in  caustic  soda  solutions.  Explain. 
Aluminum  forms  complex  anions  in  NaOH. 

Part  2.  Solubility  of  Metals  in  Salt  Solutions.  Clean  the  metal 
thoroughly,  and,  after  weighing,  set  aside  for  ten  days  in  a  test  tube 
with  10  cc.  of  the  salt  solution.  Cover  up  loosely  with  filter  paper. 
Shake  from  time  to  time.  Clean  the  test  piece  and  weigh  again. 
Record  the  time  and  note  any  change  in  the  metal. 

(a)  Copper  in  10  per  cent  NaCl,  test  alkalinity  of  filtered  solution 
at  end. 

(b)  Cadmium  in  10  per  cent  NH4NO3,  test  alkalinity  of  filtered 
solution  at  end,  and  also  test  for  nitrites. 

(c)  Iron  in  10  per  cent  sodium  tartrate.     Test  as  before. 

Reference.     Chem.  News,  90, 142  (1904). 

Part  3.     Passive  Iron. 

Discussion. 

The  passivity  of  iron  is  probably  due  to  an  adsorbed  and  stabilized 
film  of  a  higher  oxide,  the  formula  of  which  is  possibly  FeO2.  The 
oxide,  which  is  very  difficultly  soluble  in  HNO3,  is  formed  by  certain 
oxidizing  agents  such  as  HNO3,  NO2,  etc.,  or  when  iron  is  made 
anode  in  an  electrolytic  cell  through  which  a  sufficiently  high  current 
passes.  Passivity  is  removed  and  activity  is  restored  by  destruc- 
tion of  the  oxide  film.  Reducing  agents  may  destroy  the  film  or  the 
same  thing  may  be  done  by  making  a  passive  rod  cathode  with  a 
sufficiently  high  current.  Consult  the  Instructor. 

Procedure. 

(a)  Make  an  iron  rod  passive  in  concentrated  nitric  acid  (sp.  gr.  = 
1.4).     Wash  in  water  carefully  and  dip  in  dilute  HNO3  (sp.  gr.  1.2). 
What  happens? 

(b)  Having  immersed  the  rod  in  the  dilute  acid,  touch  the  rod  with 
a  fresh  (active)  iron  rod.     Explain.     Repeat,  touching  passive  rod 
with  zinc. 

(c)  Immerse  an  active  and  a  passive  rod  in  dilute  (1.2)  HNO3f 
taking  care  to  dip  the  active  rod  deeply  and  the  passive  rod  only 
slightly  beneath  the  surface  of  the  liquid.     Connect  the  two  rods  out- 
side of  the  cell  with  a  copper  wire.     What  happens? 

(d)  Repeat  experiment  (c),  having  a  large  surface  of  the  passive 
rod  and  only  a  small  surface  of  the  active  one  dipping  into  the  acid. 
To  understand  (c)  and  (d)  see  Bennett's  paper,  p.  220.     (Schonbein's 
experiments) . 

(e)  Immerse  a  passive  rod  in  dilute  acid  and  scratch  the  surface. 
Does  the  rod  become  active? 

Reference.  Bennett  and  Burnham:  Trans.  Am.  Electrochem. 
Soc.,  29,  217  (1916). 

37 


EXPERIMENTAL  GROUP  X 

REACTION  VELOCITY  AND  CATALYSIS 

This  group  of  experiments  is  designed  to  illustrate  in  a  semi- 
quantitative  manner  the  Law  of  Mass  Action  and  its  bearing  on  the 
velocity  of  chemical  change.  Simple  experiments  illustrating  cata- 
lysis are  also  included. 

Standard  References. 

Bancroft:     Papers  in  Jour.  Phys.  Chem.  (1917 — ). 

Henderson:     Catalysis  and  Its  Industrial  Applications  (1918). 

Herz:     Ahren's  Sammlung,  11,  103-145  (1906). 

Jobling:     Catalysis  and  Its  Industrial  Applications  (1916). 

Mellor:     Chemical  Statics  and  Dynamics  (1609). 

Ostwald:     Uber  Katalyse  (2nd  Ed.  1911). 

Rideal  and  Taylor:     Catalysis  in  Theory  and  Practice  (1920). 

van't  Hoff:     Lectures;    Vol.  1,  Chemical  Dynamics  (1898). 

van't  Hoff  (Evan) :     Studies  in  Chemical  Dynamics  (1896). 

Woker:     Die  Katalyse  (1915-16). 

Law  of  Mass  Action. 

The  rate  at  which  chemical  change  occurs  is  a  function  of  the 
concentration  of  each  of  the  substances  taking  part  in  the  reaction. 
The  rate  is  also  a  function  of  the  temperature  and  pressure  and  it  is 
affected  by  catalysts  and  by  various  other  influences,  such  as  light, 
electrical  and  surface  forces. 

The  law  is  illustrated  by  the  reaction  between  bromic  and  iodic 
acids 

6  HI  +  HBrO,  -^  HBr  +  3  H2O  +  3  I2, 

in  which  the  course  of  the  reaction  can  be  followed  color imetrically, 
using  starch  as  an  indicator. 

The  rate  at  which  iodine  is  set  free  is  directly  proportional  to  the 
ion  concentrations  of  iodine  and  bromate  and  to  the  square  of  the 
concentration  of  hydrogen  as  ion.  Clark:  Jour.  Phys.  Chem.,  10, 
700  (1906).  If  one  keeps  the  concentration  of  hydrogen  ions  con- 
stant and  does  not  allow  the  volume  of  the  solution  to  vary,  the 
velocity  with  which  iodine  is  liberated  at  any  moment  is  expressed  in 
terms  of  the  mass  law  by  the  equation 

_^  =  k(a  —  x)(b  —  x)  (1) 

dt 

in  which  a  and  b  refer  respectively  to  the  amount  of  iodine  and  bro- 
mate present  as  ions  at  the  beginning  of  the  experiment  and  are  there- 
fore proportional  to  the  initial  quantity  of  HI  and  HBrO3,  while  x 
refers  to  the  amount  of  iodine  or  bromate  ions  used  up  and  is  accord- 
ingly proportional  to  the  quantity  of  free  iodine  liberated. 

38 


If  the  reaction  is  allowed  to  proceed  for  a  relatively  short  time  only 
and  in  such  a  way  that  x  is  small  by  comparison  with  a  and  b,  the 
velocity  equation  takes  the  form 

^X=kab  (2)' 

whence,  on  integrating  between  the  limits  x  =  o  and  x  =  Xi ;  t  =  o, 
and  t  =  t,  the  following  expression  results : 

t  =  constant  ^L  (3) 

ab 
General  Procedure. 

In  the  experiments  which  follow  iodide  and  bromate  are  mixed  in 
acid  solution  and  the  reaction  is  allowed  to  proceed  until  a  definite 
constant  quantity  of  iodine  is  liberated,  as  determined  by  the  forma- 
tion of  a  definite  "standard"  blue  color  with  starch  as  indicator.  The 
initial  quantities  of  iodide  and  bromate  are  varied  and  the  time 
required  to  reach  the  standard  blue  is  determined  by  means  of  a  stop- 
watch. 

Under  these  experimental  conditions,  it  is  evident  from  equation 
(3)  that  the  time  required  to  reach  a  standard  blue  at  constant 
temperature  and  volume  varies  inversely  as  the  product  of  the  initial 
quantities  of  iodide  and  bromate,  as  long  as  the  amount  of  iodine  set 
free  is  small.  It  is  also  obvious  that  this  statement  becomes  less 
exact  as  the  depth  of  the  standard  blue  becomes  greater. 

For  comparison  times,  the  relative  values  of  a  and  b  may  be  sub- 
stituted for  absolute  values. 

EXPERIMENT  1 

Mass  Action 
Acid  Mixture. 

800  cc.  distilled  water 
26  cc.  N/2  HC1  (shelf) 
20  cc.  starch  solution 

To  prepare  the  starch  solution  rub  one  gram  of  starch  with  5  cc.  of 
cold  water  in  a  mortar;  pour  150  cc.  of  boiling  water  over  it,  allow 
the  undissolved  part  to  settle,  and  decant  the  supernatant  liquid. 

Standard  Blue. 

Take  two  100  cc.  bottles  (glass  stoppered)  and  in  one  make  a 
standard  blue  solution  as  follows: 
80  cc.  distilled  water 

2  cc.  starch  solution  (described  above) 
3-6  drops  "iodine  mixture"  (shelf) 

Procedure. 

Part  1 .     In  the  test  bottle  place  80  cc.  acid  mixture 

1  cc.  N/2  KBrO3  (shelf) 
1  cc.  N/2  KI  (shelf) 

in  the  order  named.  Add  the  KI  quickly  and  take  the  time  from  the 
moment  it  is  added.  Shake  at  the  moment  of  adding  KI  and  note 
the  time  required  for  the  solution  to  assume  the  same  blue  as  the 
standard.  Run  a  parallel. 

39 


Notes. 

Place  the  standard  and  the  test  bottle  against  a  white  background. 
Avoid  using  a  standard  with  too  deep  a  blue.  The  time  taken  in 
Part  1  should  not  exceed  two  minutes  nor  be  less  than  one  minute. 

Be  careful  to  work  throughout  at  constant  temperature  (20°  C.). 
Record. 

Part  2.     80  cc.  acid  mixture 

2  cc.  bromate 

1  cc.  iodide 
Shake.     Note  time  as  before.     Run  a  parallel. 

Part  3.     80  cc.  mixture  Part  4.     80  cc.  mixture 

1  cc.  bromate  2  cc.  bromate 

2  cc.  iodide  2  cc.  iodide 

Shake.     Note  time.     Run  a  Shake.     Note   time.     Run   a 

parallel.  parallel. 


EXPERIMENT  2 
Catalytic  Effect  of  Acids 

The  effect  of  acids  in  accelerating  certain  chemical  reactions  is 
roughly  proportional  to  their  electrical  conductivity.     The  effect  is 
dependent  primarily  on  the  hydrogen  ions. 
Prepare  a  mixture  as  follows : 
400  cc.  water 
10  cc.  bromate  (shelf) 
10  cc.  iodide  (shelf) 
10  cc.  starch  solution 

Part  1.  Take  80  cc.  of  the  above  mixture  in  a  100  cc.  bottle,  add 
2  cc.  N/2  HC1.  Shake.  Note  time  required.  Run  a  parallel. 

Part  2.  Take  80  cc.  of  mixture  and  2  cc.  of  N/2  H2SO4.  Shake. 
Note  time  required.  Run  a  parallel. 

Part  3.  Take  80  cc.  of  mixture  and  2  cc.  of  N/2  CH3COOH. 
Note  time  required.  Shake.  Run  a  parallel.  Explain. 

EXPERIMENT  3 
Catalytic  Effect  of  Ferrous  Sulphate 

Mixture  of  160  cc.  H2O. 

8  cc.  KI  (shelf) 
8  cc.  KBrO3  (shelf) 
4  cc.  starch  solution 

Part  1.  Take  80  cc.  of  the  above  mixture  and  10  cc.  of  N/2  acetic 
acid.  Shake.  Note  time  required  to  become  blue. 

Part  2.  Take  80  cc.  of  the  mixture  and  10  cc.  of  N/2  acetic  acid 
to  which  is  added  one  drop  of  neutral  saturated  FeSO4.  Proceed  as 
before.  Explain. 

40 


Part  3.  To  25  cc.  of  an  extremely  dilute  solution  of  chromic  acid 
(CrO3)  add  a  little  starch  solution. 

(a)  To  5  cc.  of  this  solution  add  2  to  3  drops  of  KI  solution.     Note 

time  as  before. 

(b)  To  5  cc.  of  the  solution  add  KI  as  before  and  a  little  iron  dust. 

Note  time. 

(c)  To  another  5  cc.  portion  add  KI  and  a  few  drops  of  a  ferrous 

sulphate  solution.     Note  time. 

(d)  To  another  5  cc.  portion  add  KI  and  a  few  drops  of  ferric 

sulphate  solution.     Note  time. 

Part  4.  (a)  Mix  in  the  following  order:  Dilute  CrO3  solution, 
ferrous  sulphate  solution  and  starch;  shake  and  wait  ten 
minutes;  then  add  KI.  Note  time  to  reach  standard  blue 
after  adding  KI. 

(b)  Mix  in  the  following  order:  CrO3  solution,  KI  solution  and 
starch;  wait  ten  minutes ;  then  add  ferrous  sulphate.  Note 
time  after  adding  FeSO4. 

Compare  (a)  and  (b)  and  explain. 

EXPERIMENT  4 
Hydrolysis  of  an  Ester — Catalysis 

Place  50  cc.  of  distilled  water  and  5  cc.  of  ethyl  acetate  in  a  clean, 
glass  stoppered  bottle.  Shake  thoroughly  and  titrate  duplicate 
samples  (2  cc.)  with  N/10  NaOH,  phenolphthalein  as  indicator. 

In  a  second  bottle  place  50  cc.  N/2  HC1  plus  5  cc.  ethyl  acetate. 
Shake  and  titrate  as  before. 

Set  both  bottles  aside  for  24  hours  (shaking  occasionally)  and  again 
titrate  duplicate  samples  (2  cc.).  . 

Note  differences  and  explain.  How  is  this  phenomenon  used  to 
measure  the  strength  of  acids? 

EXPERIMENT  5 
Reactions  in  Heterogeneous  Systems 

Part  1.  Size  of  Particles.  Whenever  one  of  the  reacting  sub- 
stances is  a  solid,  the  speed  of  the  reaction  is  a  function  of  the  surface 
area  of  the  solid,  or  more  accurately,  of  the  surface  per  unit  weight  of 
solid  (specific  surface).  The  specific  surface,  in  turn,  is  a  function  of 
the  size  of  the  particles  and  increases  rapidly  as  the  particles  become 
smaller.  Read  W9  Ostwald:  Grundriss  der  Kolloidchemie,  30 
(1912). 

Prepare  about  2  grams  of  finely  divided  copper  by  placing  some 
granulated  zinc  in  a  concentrated  solution  of  CuSO4.  Shake  from 
time  to  time  to  remove  the  finely  divided  copper  from  the  zinc. 
After  most  of  the  copper  has  been  precipitated,  remove  the  zinc, 
wash  the  precipitate  with  water  and  dry  in  an  air  bath.  Mix  the 
finely  divided  metal  with  powdered  sulphur  and  ignite  cautiously 
with  a  match.  What  is  formed? 

41 


Dissolve  sulphur  in  CS2  and  into  this  solution  dip  a  clean  strip  of 
copper.  What  is  the  substance  formed  on  the  copper? 

Show  how  this   experiment   illustrates  the  principle  discussed. 

Part  2.  Protecting  Films,  (a)  Clean  a  strip  of  aluminum  foil  by 
immersing  it  in  10  per  cent  NaOH.  Rinse  and  plunge  the  wet 
metal  quickly  into  clean  mercury.  Hold  it  there  until  amalgamated. 
Remove  and  rub  off  the  excess  of  mercury  adhering  to  the  aluminum, 
then  expose  to  the  air.  What  happens?  Explain. 

(b)  Place  freshly  amalgamated  aluminum  in  contact  with  warm 
water.     What  happens?     Compare  with  sodium. 

(c)  Dip  a  rod  of  metallic  magnesium  into  warm  water.     What 
happens? 

(d)  Dip  a  rod  of  metallic  magnesium  into  warm  NH4C1  solution. 
What  happens?     Explain. 

The  passivating  films  might  be  regarded  as  negative  catalysts. 

References.  Ostwald:  Prin.  Inorg.  Chem.,  560  (1900);  Wis- 
licenus:  Jour.  Praktische  Chemie,  (2),  54,  41  (1896). 

Note.  The  amalgamated  aluminum  may  be  prepared  by  cleaning 
the  metal  in  10  per  cent  NaOH,  rinsing  carefully  and  then  dipping  the 
wet  metal  into  dilute  mercuric  chloride. 

Part  3.     Reactions  between  Solids.     Incompatible  Hydrates. 

Use  small  quantities  in  proportions  approximately  equivalent. 
Weigh  out  roughly,  except  in  (a),  where  a  few  crystals  are  enough. 

(a)  Grind  together  in  a  mortar  HgCl2  +  KI. 

(b)  "  "  ""  "  Na2SO4-10H2O  +  NH4NO3. 

(c)  "  •  "  "  "  "  (NH4)2SO4  +  NaNO3. 
<d)  "  "  "  "  "  K2S04  +  NH4N03. 

(e)  "  "  ""  "  MgS04-7H2O  +  NH4N03. 

(f)  "  '"  "  "  "  CuSQ4-5H2O +  NH4NO3. 

References,  van't  Hoff:  Studies  in  Chem.  Dynamics,  173 
(1896);  Schiff:  Liebig's  Annalen,  114,  68  (1860). 

(g)  Grind  together  5  g.  NHCNS  and  10  g.  crystalline  barium 
hydroxide — Ba(OH)2  8H2O.     What  happens?     Explain. 

Part  4.  Halogen  Carriers.  Support  a  250  cc.  distilling  flask  upon 
a  ring  stand  and  connect  its  side  arm  with  a  funnel  the  mouth  of  which 
dips  just  below  the  surface  of  a  caustic  soda  solution."  Place  in  the 
flask  2  cc.  of  bromine.  Provide  a  cork  stopper  for  the  flask.  Now 
pour  into  the  flask  15  cc.  of  benzene.  Work  at  the  hoods. 

Test  for  HBr  with  ammonia  fumes.  Then  add  about  a  quarter  of 
a  gram  of  iron  powder.  Again  cautiously  test  for  HBr.  Be  re'ady 
to  stopper  the  flask  and  leave  stoppered  until  the  reaction  is  over. 


42 


EXPERIMENTAL  GROUP  XI 

SAPONIFICATION  OF  AN  ESTER 

The  experiment  which  follows  is  designed  to  demonstrate  quanti- 
tatively the  Law  of  Mass  Action  as  applied  to  the  kinetics  of  a  simple 
irreversible  reaction.  The  reaction  to  be  studied  is  a  reaction  of  the 
second  order. 

References. 

Mellor :     Chemical  Statics  and  Dynamics  (1909) . 
Warder;     Am.  Chem.  Jour.,  3  340  (1882). 
F,  270-272;  G,  246-248;  OW,  246-252,  etc. 

EXPERIMENT  1 

Saponification  of  Ethyl  Acetate 
Solutions  Required. 

A.  N/20  NaOH  (free  from  carbonates) . 

A  carbonate-free  normal  solution  of  NaOH  is  supplied  (shelf). 
From  this  prepare  a  solution  slightly  stronger  than  N/20  being  careful 
not  to  waste  any  of  the  carbonate-free  sodium  hydroxide.  Make  up 
two  liters  of  solution  and  standardize  against  an  acid  of  known  titre 
(shelf).  Finally,  dilute  until  the  solution  is  exactly  N/20  and  again 
standardize  to  make  sure  that  the  work  has  been  done  correctly. 
The  normal  titre  of  the  solution  should  not  differ  from  the  required 
value  (N/20)  by  more  than  1  per  cent.  Phenolphthalein  as  indicator. 
Save  the  residue  of  this  solution  for  use  in  Group  XIV. 

B.  N/20  HC1. 

Prepare  two  liters  and  standardize  carefully  against  N/20  NaOH. 
Protect  the  burette  containing  the  latter  by  means  of  a  soda-lime  tube. 
Phenolphthalein  as  indicator.  Save  the  residue  of  .this  solution  for 
use  in  Group  XIV. 

C.  N/20  Ethyl  Acetate. 

Ethyl  acetate  being  difficult  to  obtain  pure,  it  is  necessary  to  pre- 
pare this  solution  as  follows:  To  800  cc.  distilled  water,  contained 
in  a  liter  glass  stoppered  graduated  cylinder,  add  5  cc.  of  redistilled 
ethyl  acetate  (special  reagent).  Stopper  quickly  to  prevent  loss  of 
ester  by  volatilization,  and  shake  thoroughly  to  dissolve. 

In  a  100  cc.  glass  stoppered  bottle  place  exactly  25  cc.  of  N/20 
sodium  hydroxide  (burette)  and  to  this  add  from  a  pipette  (cali- 
brated) exactly  10  cc.  of  the  ethyl  acetate  solution.  Replace  the 
stopper  quickly  and  securely  and  heat  the  bottle  in  a  water  bath  for 
thirty  minutes,  or  until  the  ester  is  completely  saponified.  Remove 
from  bath  and  cool,  add  a  few  drops  of  phenolphthalein  and  determine 

43 


the  excess  of  sodium  hydroxide  by  titration  with  N/20  HC1.  Run  in 
duplicate.  The  solution  should  be  more  concentrated  than  N/20 
at  this  point. 

Then  calculate  the  volume  of  water  necessary  to  dilute  the  ethyl 
acetate  exactly  to  N/20,  allowing  for  the  amount  already  withdrawn. 
Saponify  as  before  in  order  to  verify  the  work.  The  ethyl  acetate 
solution  should  now  be  N/20  +_1  per  cent. 

Procedure. 

Part  1.  Adjust  the  automatic  thermostat  to  25°  C.,  or  if  this  is  not 
available  use  a  large  pan  or  beaker  of  water  kept  at  25°  +  0.1°  C. 
Measure  the  temperature  with  a  thermometer  graduated  to  tenths. 
Stir  with  compressed  air. 

Add  exactly  250  cc.  of  N/20  NaOH  to  a  500  cc.  glass  stoppered 
bottle.  Place  in  the  thermostat  and  shake  occasionally.  In  a  glass 
stoppered  measuring  flask  (250  cc.)  place  an  equal  amount  of  N/20 
ethyl  acetate.  Place  in  a  thermostat. 

When  both  solutions  have  reached  25°  C.  quickly  pour  ester  into  a 
bottle  containing  NaOH,  replace  the  stopper  and  shake  instantly. 
Start  the  stop-watch  at  the  moment  of  mixing  and  at  the  same  time 
read  the  hour  and  minute  on  a  watch,  in  case  the  stop-watch  should 
prove  faulty.  The  reaction  begins  at  the  instant  of  mixing. 

Pipette  out  10  cc.  samples  at  the  following  times: 

2,  3,  5,  8,  12,  16,  20,  25,  30,  40,  50,  60,  80,  120  minutes. 

At  the  desired  moment  stop  the  reaction  by  emptying  the  pipette 
into  an  accurately  known  volume  (about  7  cc.)  of  N/20  HC1  con- 
tained in  50  cc.  of  water  -f-  1  drop  phenolphthalein  in  an  Erlenmeyer 
flask.  Add  the  acid  from  a  burette. 

Determine  as  soon  as  possible  the  excess  of  N/20  HC1  by  titrating 
with  N/20  NaOH.  cf.  Group  X  Experiment  5. 

Shake  the  bottle  in  the  thermostat  every  two  minutes. 

Precautions. 

This  is  an  experiment  requiring  accurate  manipulation.  Burettes 
and  pipettes  should  be  calibrated  and  placed  in  cleaning  mixture 
for  at  least  twenty-four  hours  before  use.  While  in  use,  see  that  they 
are  kept  filled  with  solution  or  distilled  water,  because  drying  in  air 
causes  glassware  to  acquire  a  grease-like  film.  When  reading  bur- 
rettes  try  to  estimate  to  hundredths  of  a  cubic  centimeter. 

Computations. 

From  the  data  recorded  during  the  experiment  compute  the  num- 
ber of  cc.  of  N/20  NaOH  consumed  by  the  ester  during  each  of  the 
time  intervals.  If  this  experiment  is  carried  out  properly  these 
values  should  rise  from  zero  at  the  beginning  to  nearly  5  cc.  at  the  end. 
Draw  a  curve  between  cc.  of  NaOH  consumed  as  ordinates  and 
time  in  minutes  as  abscissae. 


44 


Using  the  equation  of  a  second  order  reaction 

k  =  _* (1) 

at  (a  —  x) 

compute  values  of  the  velocity  coefficient  k  corresponding  to  the 
different  times. 

Part  2.  Dilute  exactly  250  cc.  of  the  ethyl  acetate  solution  to  500 
cc.  making  it  N/40.  Do  the  same  with  250  cc.  of  the  NaOH.  Then 
repeat  Part  1  and  draw  a  curve  between  cc.  of  N/40  NaOH  used  up 
and  time  in  minutes. 

Compare  the  curves  obtained  in  Parts  1  and  2.  Determine  in 
each  case  the  time  required  for  one-half  of  the  original  NaOH  to 
disappear.  How  do  these  times  compare  and  how  are  they  related 
to  the  initial  concentrations  of  ester  and  base? 

From  these  results  determine  the  order  of  the  reaction. 


45 


EXPERIMENTAL  GROUP  XII 

THE  STUDY  OF  A  REACTION 

The  experiments  which  comprise  this  group  constitute  a  detailed 
study  of  the  reaction  between  oxalic  acid,  potassium  permanganate 
and  sulphuric  acid,  in  aqueous  solution: 

5  (COOH)2  +  2  KMn04  +  3  H2SO4  =  2  MnSO4  +  K2SO4  + 
10  CO2  +  8  H2O 

The  reaction  is  familiar  to  every  chemist  because  of  the  important 
part  it  plays  in  volumetric  analysis. 

Reference.     Harcourt  and  Esson:     Jour.   Chem.   Soc.,   20,  460 

(1867). 

Discussion. 

The  reaction  as  written  above  is  really  the  result  of  a  series  of 
simpler  reactions.  The  reaction  will  be  studied  by  means  of  velocity 
determinations  made  by  ascertaining  how  much  permanganate  is  used 
up  under  definite  experimental  conditions,  and  by  systematically 
varying  these  conditions. 

The  important  factors  constituting  the  experimental  conditions 
are  the  following: 

(1)  Amount  of  KMnO4. 

(2)  Amount  of  oxalic  acid. 

(3)  Amount  of  H2SO4. 

(4)  Amount  of  MnSO4. 

(5)  Amount  of  K2SO4. 

(6)  Amount  of  CO2  in  solution. 

(7)  Volume  of  solution  (amount  of  water). 

(8)  Temperature. 

(9)  Pressure. 

(10)  Illumination. 

(11)  Time. 

Factors  1  to  7  inclusive  are  concentration  factors.  In  the  present 
study  temperature,  pressure,  and  illumination  are  kept  as  nearly 
unchanged  as  possible  without  special  precautions  and  the  experi- 
ments are  carried  out  at  constant  volume.  The  work  is  done  in 
open  vessels  so  that  factor  (6)  is  practically  constant  throughout. 
Arbitrary  values  are  assigned  to  four  of  the  first  five  factors  while  the 
fifth  is  being  varied  in  a  systematic  manner.  The  time  during  which 
the  reaction  takes  place  (factor  11)  is  four  minutes. 

46 


Solutions. 

Prepare  the  following  solutions: 

(1)  Potassium  permanganate  1. 58  g.  per  liter. 

(2)  Oxalic  acid  3.15  "  "  " 

(3)  Sulphuric  acid  1.47  "  "  " 

(4)  Manganous  sulphate  2.23  "  "  " 

(5)  Potassium  sulphate  (500  cc.)  0.87  "  "  " 

(6)  Potassium  iodide  (500  cc.)  25.00  "  "  " 

(7)  Sodium  thiosulphate  2.48  "  "  " 

The  first  five  of  the  above  solutions  are  of  such  strength  in  every 
case  that  one  liter  contains  — -  of  the  number  of  molecules  taking 

£i\J\J 

part  in  the  reaction.     Thus,  in  the  case  of  the  sulphuric  acid: 
3  x  M.  W.  =  3  x  98  =  294,  and 

?^  =  1.47  (grams  per  liter) 
200 

• 

Comparison  of  Solutions. 

(a)  Titrate    the     oxalic    acid     against     the    permanganate     in 
strongly  acid  solution.     These  solutions  must  be  equivalent  i.  e.  10  cc. 
KMn04  =  10  cc.  (CO2H)2. 

(b)  Decompose  5  cc.  of  the  KMnO4  solution  by  adding  15  cc.  of 
KI  solution.     Determine  the  amount  of  iodine  liberated,  by  titrating 
with  the  thiosulphate.     Five  cc.  of  the  KMnO4  should  require  about 
25  cc.  of  the  thiosulphate. 

In  titrating,  the  iodine  with  thiosulphate,  do  not  add  the  starch 
indicator  until  most  of  the  iodine  has  been  reduced.  When  the 
solution  has  acquired  a  pale  straw  color,  add  the  starch.  A  blue 
color  should  appear.  Practice  this  titration  until  satisfactory  end- 
points  are  obtained. 

The  starch  indicator  may  be  prepared  by  rubbing  a  gram  of  arrow- 
root starch  into  a  paste  with  cold  water,  and  to  this  paste  adding 
about  200  cc.  of  boiling  water. 

Experimental  Procedure. 

The  required  amounts  of  all  of  the  reacting  substances  except  the 
permanganate,  are  mixed,  diluted  to  100  cc.  and  placed  in  Erlen- 
meyer  flasks.  These  are  then  allowed  to  come  to  the  same  tempera- 
ture. Take  the  temperature  of  each  mixture  and  record  it.  The 
permanganate  is  then  added  quickly,  the  flask  shaken  immediately 
and  the  time  taken  with  a  watch.  The  reaction  commences  with  the 
addition  of  the  permanganate. 

After  exactly  four  minutes  have  elapsed,  the  excess  of  undecom- 
posed  KMnO4  is  destroyed  by  adding  an  excess  of  the  KI  solution 
(10  to  15  cc.)  and  the  iodine  set  free  is  determined  with  thiosulphate 
solution,  using  starch  as  indicator. 

47 


The  reactions  are  as  follows : 

KMnO4  +  5  KI  +  4  H2SO4  =  3  K2SO4  +  51+4  H2O  +  MnSO4 
and 

2  Na2S203  +  2  I  =  2  Nal  +  Na2S4Ofi 

The  amount  of  decomposed  permanganate  is  proportional  to  the 
volume  of  thiosulphate  used  in  reducing  the  iodine,  and  we  may  thus 
determine  the  permanganate  used  up  in  the  reaction,  the  thiosulphate 
titre  of  the  permanganate  solution  being  known. 


Part  1.     Variation  of  Sulphuric  Acid. 


KMnO4 

H2C204 

H2S04 

(cc) 

(cc) 

(cc) 

(a)     5 

5 

5 

(b)     5 

5 

10 

(c)     5 

5 

15 

(d)     5 

5 

25 

(e)     5 

5 

40 

(f)     5 

5 

60 

Part  2. 

Variation  of  Manganous 

Sulphate. 

KMn04 

H2C204 

H2SO4 

(cc) 

(cc) 

(cc) 

(a)     5 

5 

15 

(b)     5 

5 

15 

(c)     5 

5 

15 

(d)     5 

5 

15 

(e)     5 

5 

15 

(f)      5 

5 

15 

(g)     5 

5 

15 

Part  3. 

Variation  of  Oxalic  Acid. 

KMnO4 

H2C204 

H2S04 

(cc) 

(cc) 

(cc) 

(a)     5 

1 

25 

(b)    5 

2 

25 

(c)     5 

3 

25 

(d)     5 

4 

25 

(e)     5 

5 

25 

(f)     5 

6 

25 

(g)     5 

7 

25 

(h)     5 

8 

25 

(i)      5 

9 

25 

(j)      5 

10 

25 

(k)     5 

11 

25 

(D      5 

12 

25 

(m)   5 

15 

25 

(n)     5 

20 

25 

(o)     5 

25 

25 

(p)     5 

30 

25 

(q)     5 

50 

25 

MnSO4 
(cc) 
5 
5 
5 
5 
5 
5 


MnSO4 

(cc) 

0 

1 

3 

5 

8 
10 
15 


MnSO4 
(cc) 
10 
10 
10 
10 
10 
10 
10 
10 
10 
10 
10 
10 
10 
10 
10 
10 
10 


48 


Part  4.     Variation  of  Potassium  Sulphate . 

KMnO4                  H2C2O4       H2SO4         MnSO4  K2SO4 

(cc)                          (cc)              (cc)              (cc)  (cc) 

(a)  5                            5                  15                10  5 

(b)  5                             5                  15                10  25 

(c)  5                            5                  15                10  75 

Part  5. 

To  5  cc.  of  the  KMnO4  solution  in  a  test  tube  add  10  cc.  of  the 
MnSO4  solution.  Shake  and  allow  to  settle.  What  is  the  precipi- 
tate? What  color  is  the  supernatant  liquid?  Test  it  with  litmus. 

Computations  and  Curves. 

From  your  data,  compute  (in  cc.)  the  amount  of  potassium  per- 
manganate used  up  in  the  reaction  after  four  minutes.  Call  these 
numbers  "y."  Tabulate  the  values  of  y  along  with  the  correspond- 
ing values  of  the  substance  undergoing  variation,  called  "x."  With 
values  of  x  as  abscissas  and  of  y  as  ordinates,  draw  four  curves  pictur- 
ing the  results  obtained  in  each  of  the  four  parts. 


49 


EXPERIMENTAL  GROUP  XIII 

REVERSIBLE  REACTIONS  AND  CHEMICAL  EQUILIBRIUM 

The  following  experiments  deal  particularly  with  chemical  reac- 
tions which  occur  readily  in  both  directions  and  are  therefore  dis- 
tinctly reversible,  tending  to  reach  a  condition  of  equilibrium. 
Several  examples  of  reactions  of  this  type  have  already  been  studied, 
notably  in  Groups  VIII  and  IX. 

References.     See  under  Group  X. 

EXPERIMENT  1 
Homogeneous  Chemical  Equilibrium 

This  is  illustrated  very  simply  by  the  equilibrium  between  the 
reciprocal  pairs,  ammonium  thiocyanate-ferric  chloride  and  ferric 
thiocyanate-ammonium  chloride.  The  amount  of  ferric  thiocyanate 
formed  in  solution  may  be  estimated  by  the  intense  red-brown  color 
that  the  undissociated  salt  imparts. 

"If  the  reaction  is  represented  by 

3  NH4CNS  +  FeCl3  =  Fe  (CNS)3  +  3NH4C1 

and  the  amount  of  ferric  sulphocyanate  is  judged  by  the  depth  of 
color  of  the  solutions,  the  reaction  between  equivalent  quantities 
must  be  regarded  as  incomplete." 

Procedure. 

The  following  solutions  will  be  found  as  shelf  reagents : 

Solution  A.     Ammonium  thiocyanate  38  g.  NH4CNS  per  liter. 

Solution  B .     A  mixture  of  the  following : 

Ferric  chloride  30  g. 

Concentrated  hydrochloric  acid  115  cc. 

Water,  1000  cc. 

Take  equal  volumes  of  solutions  A  and  B,  5  cc.  of  each.  Dilute  to 
2  liters.  Stir  thoroughly. 

The  solution  should  be  a  definite  orange  in  color. 
Divide  into  five  400  cc.  portions. 

To  the  is  added  Color  becomes 

First  portion  5  cc.  NH4CNS  solution  ? 

Second  portion  5  cc.  FeCl3  solution  ? 

Third  portion  50  cc.  sat.  NH4C1  solution  ? 

Fourth  portion  solution  of  HgCl2  ? 

The  first  portion  is  kept  for  comparison.     Explain  all  results. 

Reference.  Miller  and  Kenrick:  Jour.  Am.  Chem.  Soc.  22,  292 
(1900). 

50 


EXPERIMENT  2 

Heterogeneous  Chemical  Equilibrium 
BaSO4  +  Na2CO3  =  BaCO3  +  Na2SO4 

Part  1.  In  an  evaporating  dish  over  a  water  bath  heat  together 
1/100  molecular  weight  of  BaSO4,  1/6  molecular  weight  of  Na2CO3, 
and  100  cc.  of  water.  Stir  constantly  and  replace  water  that 
evaporates. 

After  heating  for  one  hour,  test  the  supernatant  liquid  for  sul- 
phates, taking  care  to  decompose  the  carbonates  before  testing. 

Wash  the  residue  by  decantation  and  finally  on  a  filter  until  the 
wash  water  gives  no  test  for  carbonates. 

After  washing  to  free  from  soluble  carbonates  test  the  residue. 
What  is  it?  Explain. 

Part  2.  Take  1/100  molecular  weight  of  BaCO3,  add  100  cc.  of 
water,  then  add  1/6  molecular  weight  of  Na2SO4  .  10  H2O. 

Follow  the  same  procedure  as  in  Part  1. 

Test  the  supernatant  liquid  for  carbonates. 

Test  the  residue,  after  washing,  with  HC1.  Is  there  a  residue  after 
treating  with  HC1?  What  is  this  residue? 

Explain. 

Reference.  Walker,  290  (1913);  Mellor:  Chem.  Statics  and 
Dynamics,  243  (1909). 


EXPERIMENT  3 
Heterogeneous  Chemical  Equilibrium 

.     2  NaCl  +  H2SO4  =  2  HC1  +  Na2SO4 
Part  1.     Concentrated  H2SO4  is  poured  into  its  own  volume  of  a 

saturated  solution  of  sodium  chloride  in  a  small  evaporating  dish. 

Warm  very  gently  in  the  hood.     Set  aside  until  crystallization  begins, 

then  pour  the  liquid  off  and  dry  the  crystals  on  a  piece  of  unglazed 

porcelain. 

The  product  is  sodium  sulphate  containing  hydrochloric  acid.     To 

test  for  the  former  it  is  necessary  to  get  rid  of  the  latter.     Dissolve  in 

the  least  possible  amount  of  water  and  precipitate  the  sulphate  by 

adding  absolute  alcohol. 

Filter,  and  after  drying  the  residue,  test  for  sodium  chloride  and 

sulphate. 

Reference.     Miller  and  Kenrick:     loc.  cit.  (Cf.  Expt.  1). 

Part  2.  Cover  a  crystal  of  Na2SO4  .  10  H2O  on  a  watch  glass  with 
concentrated  HC1.  After  the  reaction  is  complete  pour  off  the  acid 
on  an  unglazed  porcelain  plate,  as  before. 

To  analyze  the  resulting  product  warm  gently  in  a  test  tube  and 
remove  any  HC1  fumes  from  the  tube  by  blowing  out  with  air.  Then 
dissolve  in  water  and  test  for  sodium  chloride  and  sulphate. 

51 


EXPERIMENT  4 
Distribution  of  a  Base  between  Two  Acids 

Weigh  out  5  grams  of  Ba(OH)2  •  8  H2O  and  dissolve  in  50  cc.  of 
water.  Make  up  a  mixed  solution  of  H2SO4  and  HC1  (obtained  by 
calculation  and  reference  to  acid  tables)  containing  just  enough  of 
each  acid  to  neutralize  all  the  Ba(OH)2  in  the  first  solution.  Dilute 
this  solution  to  50  cc.  Then  mix  the  two  solutions.  Shake  well. 
After  settling,  test  the  supernatant  liquid  for  barium.  How  does  the 
base  distribute  itself  between  the  competing  acids?  Why  is  H2SO4, 
the  "weaker"  acid,  more  active  in  this  case?  Define  the  term 
"weaker"  acid. 

Explain  your  results. 

EXPERIMENT  5 
Addition  of  a  Common  Ion 
Discussion. 

The  dissociation  of  a  weakly  ionized  acid  or  base  is  greatly  reduced 
by  the  addition  of  one  of  its  neutral  salts.  According  to  the  Mass 
Law,  the  product  of  the  concentrations  of  the  two  ions  of  the  acid  is 
proportional  to  the  concentration  of  its  undissociated  portion  and 
since  the  concentration  of  the  anion  is  greatly  increased  by  the  addi- 
tion of  the  neutral  salt,  the  ratio  of  the  concentration  of  the  H  ion  to 
that  of  the  undissociated  acid  must  decrease  in  the  same  proportion. 
In  the  following  experiment,  in  order  to  show  the  difference  between 
the  concentration  of  the  hydrogen  ion  in  the  two  cases,  use  is  made  of 
the  relative  effect  of  the  acid,  in  the  absence  and  presence  of  its 
neutral  salt,  in  accelerating  the  bromate-iodide  reaction. 

Procedure. 

Make  a  standard  blue  solution.     See  X,  Experiment  1. 
Make  a  solution  as  follows:     175  cc.  of  water,  5  cc.  of  N/2  KBrO3, 
5  cc.  of  N/2  KI,  and  3  cc.  of  starch  solution. 

Part  1.     Solution  (1)     Take  90  cc.  of  the  above  mixture. 
Solution  (2)     Then  25  cc.  of  N/2  acetic  acid  and  mix  with  25  cc. 
water. 

Mix  (1)  and  (2)  and  note  time  to  reach  the  standard  blue. 

Part  2.     Solution  (1)     Take  90  cc.  of  the  above  mixture. 
Solution  (2)      Then  25  cc.  of  N/2  acetic  acid  and  25  cc.  of  N/2 
sodium  acetate  mixed. 

Mix  (1)  and  (2)  and  note  time  as  before. 


52 


EXPERIMENTAL  GROUP  XIV 

INDICATORS 

This  group  of  experiments  is  divided  into  two  parts,  the  first  com- 
prising a  study  of  several  of  the  more  common  indicators  employed 
in  acidimetry  and  alkalimetry,  particularly  methyl  orange  and  phe- 
nolphthalein.  The  second  part  comprises  the  rough  determination  of 
hydrogen  ion  concentration  by  the  use  of  a  set  of  indicators. 

References. 

Glaser:     Die  Indikatoren  (1901). 
Noyes:     Jour.  Am.  Chem.  Soc.,  32,  816  (1910). 
Prideaux:     Theory  and  Use  of  Indicators  (1917). 
Thiel:     Der  Stand  der  Indikatorenfrage,  Ahren's  Sammlung  16, 
307-422  (1911). 

SUB-GROUP  I 

STUDY  OF  INDICATORS 
Procedure. 

Prepare  the  following  solutions : 

N/20  HC1  1  liter. 

N/20  Acetic  acid  500  cc. 

N/20  NaOH  1  liter. 

N/20  NH4OH  500  cc. 

The  bases  should  be  free  from  carbonates.  Cf.  Group  XI. 
Use  calibrated  burettes  and  make  sure  that  these  are  absolutely 
clean.  Cf.  Group  I.  Before  taking  readings  allow  burettes  to  drain 
exactly  two  minutes  and  use  every  precaution  in  titrating.  Protect 
NaOH  from  CO2  in  the  air.  Never  leave  the  burettes  standing 
partly  empty  exposed  to  the  air,  but  keep  them  filled  with  distilled 
water"  when  not  in  constant  use. 

Always  use  the  same  amount  of  indicator  each  time. 
Prepare  a  standard  comparison  end-point  for  use  with  each  indica- 
tor, and  match  this  shade  and  color  carefully  each  time. 
Keep  the  temperature  as  constant  as  possible. 

EXPERIMENT  1 
Comparison  of  Indicators 

Part  1.  Titrate  10  cc.  HC1  with  NaOH.  Dilute  acid  to  50  cc. 
each  time. 

(a)  Phenolphthalein  (Ppn)  in  acid. 

(b)  Methyl  orange  (MO)  in  acid. 

(c)  Purified  litmus  (special  reagent)  in  acid. 

53 


Note.  Acid  and  base  should  be  very  closely  equivalent  with 
litmus.  Explain  different  readings  obtained  in  (a),  (b),  and  (c). 
Cf .  Experiment  4,  this  group. 

Part  2.     Titrate  10  cc.  of  HC1  with  NH4OH.     Dilute  to  50  cc. 

(a)  MO  in  acid. 

(b)  Ppn  in  acid. 

Part  3.     Titrate  10  cc.  acetic  acid  with  NaOH.     Dilute  to  50  cc. 

(a)  Ppn  in  acid. 

(b)  MO  in  acid. 

Part  4.     Titrate  10  cc.  acetic  acid  with  NH4OH.       Dilute  to  50  cc. 

(a)  MO  in  acid. 

(b)  Ppn  in  acid. 

From  your  results  draw  conclusions  as  to  the  proper  indicator  to 
use  under  the  various  conditions.  MO  as  indicator  seems  to  behave 
as  a  weak  base;  Ppn,  as  a  weak  acid.  Cf.  Waddell:  Jour.  Phys. 
Chem.,2,171  (1898). 

EXPERIMENT  2 
Indicators  as  Acids  or  Bases 

Indicators  are -weak  acids  or  weak  bases.  Is  there  therefore  any 
difference  in  the  amount  of  acid  or  base  required  for  neutralization, 
depending  on  whether  the  indicator  is  placed  in  the  acid  or  in  the 
base? 

Part  1.  Titrate  the  number  of  cc.  of  NaOH  required  to  neutralize 
the  acid  in  Experiment  1,  Part  la,  with  HC1,  adding  the  indicator  to 
the  base.  Dilute  to  50  cc. 

Part  2.  Titrate  the  number  of  cc.  of  NaOH  required  to  neutralize 
the  acid  in  Experiment  1,  Part  Ib,  with  HC1,  adding  the  indicator  to 
the  base.  Dilute  to  50  cc. 

EXPERIMENT  3 
Effect  of  Heat  on  Indicators 

Part  1.  Take  10  cc.  of  HC1,  dilute  to  50  cc.,  add  Ppn  and  nearly 
neutralize  with  NaOH.  Then  heat  to  80-90°  and  complete  the 
titration  at  this  temperature. 

Part  2.  Take  10  cc.  of  HC1,  dilute  to  50  cc.,  add  MO  and  nearly 
neutralize  with  NaOH.  Heat  to  70-80°  and  complete  the  titration 
at  this  temperature. 

EXPERIMENT  4 
Effect  of  Volume 

The  neutral  (end-point)  color  of  an  indicator  occurs  at  a  definite 
concentration  of  hydrogen  ions  in  the  solution.  Study  the  table  in 
Washburn  333  and  posted  in  the  laboratory.  The  hydrogen  ion 
concentration  of  the  end-point  is  different  for  the  different  indicators. 
With  this  in  mind  and  remembering  that  concentration  is  defined  as 
mass  divided  by  volume,  perform  the  following : 

54 


Part  1.     Take  10  cc.  of  HC1,  dilute  to  250  cc.,  add  Ppn  and  titrate 
with  NaOH. 

Part  2.     Take  10  cc..of  HC1,  dilute  to  500  cc.,  add  Ppn  and  titrate 
with  NaOH. 

Part  3.     Take  10  cc.  of  HC1,  dilute  to  250  cc.,  add  MO  and  titrate 
with  NaOH. 

Part  4.     Take  10  cc.  of  HC1,  dilute  to  500  cc.,  add  MO  and  titrate 
with  NaOH. 

EXPERIMENT  5 

Phosphoric  Acid 
Discussion. 

Phosphoric  acid  dissociates  in  three  stages: 
(1)  H,PO4  =H+  +  H2P04 


(2)  H2PO    =  H+  +  HPO     (somewhat) 

(3)  HPO7  =  H+  +  PO4  =  (very  slightly) 

Accordingly  phosphoric  acid  is  really  a  fairly  strong  monobasic 
acid,  but  as  a  dibasic  acid  it  is  weak. 

On  adding  NaOH,  the  reaction  first  takes  the  following  course: 

H3PO4  +  NaOH  =  NaH2PO4  +  H2O 

The  ions  are  Na  +  and  H2PO4—  .  Referring  to  stage  2  in  the  ioniza- 
tion  of  phosphoric  acid  it  is  seen  that  H2PO4  —  also  ionizes  somewhat 
into  H+  and  HPO4  =  .  The  hydrogen  ions  are  so  few,  however,  that 
their  concentration  is  not  sufficient  to  turn  MO  red,  but  is  sufficient  to 
render  Ppn  colorless.  On  adding  a  second  molecule  of  NaOH,  the 
reaction  becomes: 

Na2HPO4  +  NaOH  =  Na2HPO4  +  H2O 

The  ions  are  now  Na+  and  HPO4=.  Since  HPO4=  gives  scarcely 
any  H  +  and  PO3  =  ions  (stage  3)  and  does  not  react  readily 
with  NaOH,  a  very  small  quantity  of  base  in  excess  of  two  equivalents 
will  give  a  solution  sufficiently  alkaline  to  turn  Ppn  pink.  Read 
Stieglitz,  I  103.  Do  your  results  check  with  the  theory? 

Procedure. 

Part  1.     Titrate  10  cc.  of  M/20  phosphoric  acid    (shelf)   with 
NaOH,  as  indicator. 

Part  2.     Titrate  10  cc.  of  M/20  phosphoric  acid  with  NaOH. 
Ppn  as  indicator.     Walker,  359  (1913). 

EXPERIMENT  6 

Titration  of  Carbonates.     Effect  of  CO2 
Dissolve  0.3  gram  of  NaiCO3  in  120  cc.  of  H2O 

Part  1.     Take  20  cc.  of  this  solution  and  titrate  with  N/20  HC1 
(MO)  as  indicator.     When  the  end-point  is  reached,  heat  to  boiling. 

55 


Part  2.  Take  20  cc.  of  this  solution,  heat  to  boiling,  and  titrate 
with  N/20  HC1  (MO)  as  indicator. 

Part  3.  Take  20  cc.  of  this  solution  and  titrate  with  N/20  HC1 
with  Ppn  as  indicator.  When  the  end-point  is  reached,  heat  to 
boiling. 

Part  4.  Take  20  cc.  of  this  solution,  heat  to  boiling,  and  titrate 
with  N/20  HC1  with  Ppn  as  indicator. 

Part  5.     To  a  solution  of  Na2CO3  add  Ppn. 
To  a  solution  of  Na2CO3  add  MO. 
To  a  solution  of  NaHCO3  add  Ppn. 
To  a  solution  of  NaHCO3  add  MO. 

Part  6.  To  a  dilute  solution  of  NaOH  add  Ppn.  Pass  CO2  into 
the  solution.  Does  the  pink  color  disappear?  Does  it  reappear  on 
passing  in  more  CO2? 

Repeat,  using  MO.     Explain  all  results. 

Part  7.  To  a  solution  of  Na2CO3  add  Ppn,  then  pass  CO2into  the 
solution. 

Repeat,  using  MO.     Explain. 

Hint.     Consider  the  reaction  as  occurring  in  two  stages: 
2  NaOH  +  C02  =  Na2CO3  +  H2O 
Na2C03  +  C02  +  H20  =  2  NaHCO3 
Compare  with  phosphoric  acid  in  Experiment  5  above.     Explain. 

EXPERIMENT  7 
Miscellaneous 

Part  1.  To  20  cc.  of  alcohol  plus  a  few  drops  of  phenolphthalein 
add  several  drops  of  aqueous  ammonia,  and  shake  the  solution. 
Water  is  added  slowly  up  to  5  cc.  Then  add  25  cc.  of  alcohol. 
Explain. 

References.     Elements  of  Phys.  Chem.,  295  (1907). 
Hildebrand's  explanation,  Jour.  Am.  Chem.  Soc.,  30,  1914  (1908). 
Jones'  explanation,  Am.  Chem.  Jour.,  18,377  (1896). 

Part  2.  Add  a  little  Ppn  to  concentrated  H2SO4.  Add  a  little 
Ppn  to  concentrated  NaOH. 

Reference.     McCoy:     Am.  Chem.  Jour.,  31,  516  (1904). 

Part  3.  Divide  a  dilute  acetic  acid  solution  into  two  portions,  and 
add  MO  to  each.  To  one  add  sodium  acetate.  Show  that  this  solu- 
tion is  still  acid  to  litmus.  Explain.  Cf.  Stieglitz,  I,  113. 

SUB-GROUP  2 

HYDROGEN  ION  CONCENTRATION 
Discussion. 

Read  Prideaux  on  Indicators,  or  the  more  recent  general  texts,  such 
as  Washburn  or  Lewis.  All  aqueous  solutions,  whether  acid,  neutral 

•   56 


or  alkaline,  contain  both  hydrogen  and  hydroxyl  ions,  the  product  of 
their  concentrations  being  roughly  1.0  x  10-14  at  25°  C.  In  neutral 
solutions  these  concentrations  are  equal  and  lie  close  to  10~7gram  ions 
per  liter.  A  solution  normal  with  respect  to  hydrogen  ions  would 
represent  a  hydrogen  ion  concentration  of  10°  or  unity;  a  tenth- 
normal  solution  a  hydrogen  ion  concentration  of  10"1  or  1/10  and  so 
on ;  a  solution  normal  with  respect  to  hydroxyl  ions  would  represent 
a  hydrogen  ion  concentration  of  10~14.  It  follows  therefore  that 

The  degree  of  acidity  or  alkalinity  of  any  solution  may  be  expressed 
in  terms  of  its  hydrogen  ion  concentration. 

Sorensen  has  suggested  that  the  hydrogen  ion  concentration  be 
represented  in  terms  of  an  index  represented  by  the  symbol  PJJ« 
This  "Index"  is  the  common  logarithm  of  the  hydrogen  ion  concentration 
with  the  minus  sign  omitted.  Thus  if  PH  =  1,  the  solution  has  a 
hydrogen  ion  concentration  of  10-1  and  is  tenth-normal;  PH  =  7 
would  represent  a  neutral  solution,  and  so  on. 

When  PH  is  greater  than  7,  the  solution  is  alkaline;  when  it  is  less 
than  7,  the  solution  is  acid,  provided  one  is  dealing  with  so-called 
"room"  temperatures  (18-25°  C.). 

The  most  accurate  and  reliable  method  of  measuring  the  hydrogen 
ion  concentration  of  a  certain  solution  is  an  electrical  one  employing 
a  hydrogen  electrode.  This  electrometric  method  is  studied  in  the 
laboratory  course  in  electrochemistry,  Course  56b.  Since,  however, 
indicators  undergo  their  characteristic  color  changes  and  show  their 
neutral  colors  at  very  definite  hydrogen  ion  concentrations,  a  set  of 
indicators  may  be  used  to  measure  hydrogen  ion  concentration,  pro- 
vided the  critical  or  neutral  color  concentration  is  known  for  each 
indicator  and  the  range  covered  is  sufficiently  great. 

Cf.  Clark:     The  Determination  of  Hydrogen  Ions  (1920). 

The  indicator  method  may  be  carried  out  by  comparing  the 
unknown  solution  with  a  set  of  standard  solutions  of  known  hydrogen 
ion  concentration  and  determining  with  which  of  these  standard  solu- 
tions the  unknown  is  most  nearly  identical.  The  following  standard 
solutions  are  available  (special  reagents) : 

Standard  Solutions  of  Known  Hydrogen  Ion  Concentration 
Reference.     Noyes:     Jour.  Am.  Chem.  Soc.,  32,  822  (1910). 

(1)  PH  =  3.     Mix  570  cc.  of  N/10  acetic  acid  with  430  cc.  of 
water.     Acetic  acid  contains  6  g.  per  liter. 

(2)  PH  =  4.     Dissolve  2.7  g.  CH3COONa-3H2O  in  1  liter  of  N/10 
acetic  acid.     CH3COONa.3H2O  is  crystalline  sodium  acetate. 

(3)  PH  =  5.     Dissolve  15.0  g.  of  CH3COONa-3H,O  in  500  cc.  of 
water,  and  add  500  cc.  N/10  acetic  acid. 

(4)  PH  =  6  to  11.     Make  up  a  tenth  molecular  solution  of  Na2 
HPO4'12H2O.     Prepare  also  N/10  HC1  and  N/10  NaOH  (free  from 
carbonate).     Mix  the  solutions  as  follows: 

57 


PH 

Mix 

6 
7 
8 
9 
10 
11 

600  cc. 
700  cc. 
1000  cc. 
1000  cc. 
1000  cc. 
1000  cc. 

N/10  NaH 
N/10 
N/10 
N/10 
N/10 
N/10 

:P04  +  500  cc. 
+  3,50  cc. 
+    47  cc. 
+      5  cc. 
+  3.6  cc. 
+    36  cc. 

N/10  HC1. 
N/10     " 

N/10     " 
N/10     " 
N/10  NaOH. 
N/10  NaOH. 

These  solutions  will  be  found  on  the  reagent  shelf. 
For  other  mixtures  giving  solutions  of  known  hydrogen  ion  concen- 
tration consult  Walpole:     Biochemical  Journal  5,  207  (1911). 

Indicator  Solutions. 

0.5  per  cent  thymolphthalein  (Tpn)  in  alcohol. 

0.5  per  cent  phenol phthalein  (Ppn)  in  alcohol. 

0.5  per  cent  rosolic  acid  (RA)  in  50  per  cent  alcohol. 

0.1  per  cent  methyl  red  (MR)  in  water. 

0.1  per  cent  methyl  orange  (MO)  in  water. 

0.1  per  cent  purified  litmus  (L)  in  water. 

Extract  of  cochineal  (Coc)  in  water. 

EXPERIMENT  1 

To  Determine  Hydrogen  Ion  Concentration  Corresponding  to  Neutral 
or  Critical  Color  of  Indicators 

Noyes:     Jour.  Am.  Chem.  Soc.,  32,  824  (1910). 

Obtain  twenty-seven  test  tubes,  clean  and  dry,  then  place  in  nine 
groups  of  three.  To  each  test  tube  add  10  cc.  of  the  various  standard 
solutions  of  known  hydrogen  ion  concentration  and  to  these  0.1  cc. 
(two  drops)  of  the  various  indicators,  according  to  the  following 
scheme : 
Standard  Solutions  PH  =3  45  67  8  9  10  11 

(1)  —   MR  MR  MR  MR  Tpn  Tpn  Tpn  Tpn 

(2)  Coc  Coc  Coc  Coc  —     Ppn  Ppn  Ppn  Ppn 

(3)  MO  MO  MO  RA  RA  RA     RA 

Determine  where  the  critical  color  change  occurs.     How  do  your 
results  agree  with  what  Noyes  found,  or  with  the  table  given  in 
Washburn? 

Note  that  the  experiment,  performed  as  outlined  above,  is  only 
roughly  quantitative.  For  accurate  work  the  color  changes  should 
be  observed  in  a  colorimeter.  Washburn,  332. 

EXPERIMENT  2 
To  Determine  Hydrogen  Ion  Concentration  of  an  Unknown 

Determine  the  approximate  hydrogen  ion  concentration  of  N/10 
methylamine  hydrochloride.  Employ  the  set  of  indicators  listed 
above. 

From  your  results  compute  the  per  cent  hydrolysis  of  N/10 
methylamine  hydrochloride. 


I—  1 
1 

i 

0 

rH 

0 

O 
1—  1 

§ 

0) 

PQ 

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pS 

"o 

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OJ 

^o 

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O 

1 

l-a 

5=1 

> 

I 

OH 

K^S 

TJ 

0 

o 

1 

o 

0 

'S 

<D 

5 

g 

^0 

'o 

o 

(!) 

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<U 

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^~* 

ry^ 

u 

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2  . 

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<u 

G 

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0) 
IH 

S 

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£ 

05 

-g 

o 

True  Acidity 
(H  +  concentratic 

True  alkalinity 
(OH~concentrati 

Dimethylamido- 
azobenzene 

0) 

I 

1 

Sodium  Alizarin- 
Sulphonate 

Rosolic  Acid 

Guiac  Tincture 

Phenolphthalein 

Thymolphthalein 

Methyl  Red 

Litmus 

Cochineal 

EXPERIMENTAL  GROUP  XV 

EQUILIBRIUM  AND  THE  PHASE  RULE 

The  series  of  experiments  outlined  in  this  group  constitutes  a  study 
of  physical  and  chemical  equilibrium  from  the  point  of  view  of  the 
Phase  Rule.  It  includes  phase  equilibria  in  systems  of  one,  two  and 
three  components.  Compare  the  experiments  on  distillation  (Group 
VIII)  and  vapor  pressure  (Group  IV).  In  carrying  out  the  experi- 
mental work  keep  the  Phase  Rule  in  mind. 

References : 

Bancroft :     The  Phase  Rule  (1897) . 

Desch:     Metallography  (1913) ;   Intermetallic  Compounds  (1914). 

Duhem  (Burgess) :  Thermodynamics  and  Chemistry  (1903). 

Findlay:     The  Phase  Rule  (1917).     (FPR). 

Roozeboom :     Die  heterogenen  Gleichgewichte  (1901-1913) . 

Tammann:     Kristallizieren.und  Schmslzen  (1903). 

SUB-GROUP  1 

INVERSION  POINTS 
Discussion. 

Read  carefully  the  appendix  in  FPR,  335  (1911)  or  F,  307-315. 

Do  not  begin  experimental  work  until  you  are  thoroughly  familiar 
with  the  principles  involved.  Note  "suspended  transformation," 
FPR,  69  (1911). 

"It  frequently  happens  that  in  place  of  determining  the  complete 
concentration-temperature  curve  and  from  the  break  determining 
both  the  concentration  and  temperature  at  the  inversion  point,  one 
prefers  to'  measure  the  temperature  at  which  such  changes  occur. 
Since  a  change  in  the  solid  phase  brings  a  change  in  practically  all  the 
physical  properties,  the  close  observation  of  the  variations  of  any 
one  of  these  with  the  temperature  will  decide  at  which  temperature 
the  inversion  takes  place.  The  different  properties  whose  variations 
are  accessible  to  easy  measurement  are  crystal  form,  volume,  color, 
vapor  pressure,  conductivity,  and  electromotive  force.  The 
variation  of  the  physical  properties  is  accompanied  by  a  variation  of 
the  energy  content  so  that  by  measurement  of  the  variation  of  some 
energy  quantity  with  the  temperature,  the  inversion  point  may 
readily  be  found  by  all  the  methods;  as  in  analysis,  every  particular 
case  shows  one  method  which  ought  to  be  employed  in  preference  to 
the  others,  because  of  its  sharpness  in  detecting  the  change. 

"In  practically  all  cases  where  phase  changes  (inversions)  occur, 
there  is  a  lag  or  reluctance  to  change,  which  may  be  more  marked  in 
one  direction  than  in  the  other.  This  reluctance  to  change  gives  rise 
to  metastable  phases  and  to  metastable  equilibria.  Even  when  the 

60 


change  of  phase  (inversion)  is  actually  occurring,  time  is  required  for 
the  change  and  this  may,  and  usually  will,  introduce  a  complicating 
factor  in  the  experimental  determination  of  inversion  temperature." 

EXPERIMENT  1 
Optical  Method 

Part  1.  Determine  by  means  of  color  change  the  inversion  tem- 
perature of  mercuric  iodide.  One  component. 

Carry  out  this  determination  with  the  aid  of  a  Thiele  bulb,  as  you 
would  make  a  determination  of  the  melting  point.  Use  H2SO4  and 
heat  very  slowly.  Note  the  point  at  which  the  color  change  occurs 
with  both  rising  and  falling  temperature.  What  is  the  cause  of  the 
difference?  What  is  this  phenomenon  called? 

In  a  test  tube  heat  the  red  HgI2  until  it  becomes  yellow.     Pour 
melted  vaseline  over  some  of  the  yellow  HgI2  and  cool  quickly.     Like- 
wise cool  the  remainder  of  the  yellow  iodide  exposed  to  the  air.     Is . 
the  stability  of  the  yellow  form  affected  by  the  presence  of  vaseline? 

Reference.     FPR,  75  (1917). 

Part  2.  Following  the  same  procedure,  determine  the  inversion 
point  of  copper  potassium  chloride.  Pick  out  blue  crystals  of  the 
hydrated  double  salt  in  preference  to  the  green  ones.  Three  com- 
ponents: CuCla  •  2KC1  •  2H2O  =  CuCl2  •  KC1  +  KC1  +  2H2O. 

How  many  phases  are  in  equilibrium  at  the  inversion  point? 

EXPERIMENT  2 

Thermometric  Method  (Cooling  Curves) 
Discussion. 

If  a  system  of  phases  is  at  a  temperature  different  from  the  sur- 
roundings it  will  either  absorb  or  give  off  heat  according  to  its  tem- 
perature. If  at  any  temperature  there  occurs  in  the  system  some 
change  where  heat  is  evolved  or  absorbed  there  must  necessarily  be 
a  break  in  the  curve  of  heating  or  cooling.  Since  the  appearance  or 
disappearance  of  a  phase  is  always  accompanied  by  a  heat  change,  one 
may  easily  and  rapidly  make  the  determination  by  observing  the 
temperature-time  curve  indicating  the  rapidity  of  heating  or  cooling 
of  the  system. 

Procedure. 

Part  1.  Determine  the  inversion  temperature  of  sodium  sulphate 
decahydrate  (Glauber's  salt)  by  the  thermometric  method. 

In  a  test  tube  place  sufficient  powdered  salt  to  cover  completely  the 
bulb  of  a  large  thermometer  graduated  in  tenths.  The  test  tube 
should  be  half  full.  Place  the  test  tube  in  a  water  bath  and  beginning 
at  28°  heat  slowly  to  36°,  stirring  the  contents"  of  the  test  tube  con- 
stantly with  the  thermometer.  Raise  the  temperature  of  the  bath 
at  a  uniform  rate,  not  exceeding  one  degree  in  five  minutes. 

Read  the  temperature  on  the  thermometer  immersed  in  the  salt 
at  regular  intervals  of  two  minutes.  At  the  same  time  record  any 
changes  which  may  be  visible  in  the  contents  of  the  tube.  Draw  a 
curve  between  temperature  and  time  and  note  the  "break"  at  the 
inversion  temperature. 

61 


Next  cool  the  test  tube  and  contents  from  36°  to  28°  proceeding  as 
you  did  before.  Draw  a  cooling  curve  between  temperature  and 
time. 

If  undercooling  becomes  excessive  and  persists,  add  a  crystal 
of  decahydrate  and  stir  vigorously.  Account  for  the  sudden  rise  of 
temperature. 

How  many  components  and  phases  are  there  at  the  inversion 
point?  How  does  the  inversion  point  differ  in  this  case  from  a  melt- 
ing point?  Has  Glauber's  salt  a  melting  point? 

Part  2.  Determine  the  inversion  temperature  of  mercuric  chloride 
methylalcoholate,  HgCl2  •  CH3OH.  Saturate  methylalcohol  at  45°  C. 
with  HgCl2.  Cool  and  determine  the  temperature  at  which  HgCl2 
ceases  to  be  deposited  and  the  alcoholate  makes  its  appearance. 
The  reaction  may  be  written 

HgCl2  +  CH3OH  =  HgCl2  •  CH3OH. 

Reference.    Jour.  Phys.  Chem.,  1,  298  (1896). 

Caution.     Work  at  the  hoods. 

EXPERIMENT  3 
Dilatometric  Method  (Volume  Changes) 

The  powdered  solid  is  introduced  into  the  bulb  of  a  glass  dilato- 
meter  through  the  larger  tube  below  the  bulb.  The  capillary  tube 
is  closed  by  means  of  a  small  piece  of  glass  to  prevent  the  solid  sub- 
stance from  clogging  the  capillary.  This  piece  of  glass  may  best  be 
made  by  drawing  out  a  glass  rod,  then  forming  a  bead  at  one  end 
by  holding  it  in  the  flame  for  an  instant.  The  bulb  is  then  nearly 
filled  with  the  solid  and  the  larger  tube  sealed  off. 

The  dilatometer  must  now  be  filled  with  some  measuring  liquid, 
e.  g.,  petroleum  or  xylene.  This  is  best  done  by  attaching  an  adapter 
to  the  end  of  the  capillary  tube  by  means  of  a  rubber  stopper  fitting 
the  wide  end  of  the  adapter  and  then  connecting  the  latter  to  a  suc- 
tion pump  after  filling  with  xylene.  The  air  from  the  dilatometer 
bubbles  through  the  oil,  which,  when  the  pressure  is  released,  is 
drawn  back  into  the  dilatometer,  Cf.  F,  312  (1917). 

This  operation  is  repeated  until  all  the  air  is  withdrawn  from  the 
dilatometer  and  replaced  by  xylene.  This  capillary  tube  of  the  dila- 
tometer should  be  tapped  frequently  to  loosen  any  adhering  air 
bubbles.  Any  excess  of  xylene  may  be  removed  from  the  capillary 
by  means  of  a  long  finely  drawn  out  capillary  tube,  so  that  when  the 
dilatometer  is  placed  in  the  water  bath  the  xylene  meniscus  may 
remain  on  the  scale.  The  capillary  tube  is  not  sealed.  A  suitable 
millimeter  scale  is  used  for  reading  the  change  in  volume.  This 
method  is  especially  useful  for  determining  inversion  points  when  the 
amount  of  substance  obtainable  is  relatively  small. 

By  means  of  the  method  described  find  the  inversion  temperature 
of  sodium  thiosulphate  pentahydrate,  Na2S2O3  •  5  H2O. 

After  the  dilatometer  has  been  filled,  place  it  in  a  large  beaker  of 
water  and  starting  at  46°,  heat  to  52°  at  the  rate  of  1°  every  five 
minutes,  noting  the  change  in  volume.  Then  allow  the  dilatometer 
to  cool  very  slowly,  taking  readings  of  temperature  and  volume. 

62 


Finally,  start  at  a  temperature  about  two  degrees  below  the  inver- 
sion temperature  and  heat  to  a  temperature  of  about  two  degrees 
above,  at  a  rate  of  1°  every  ten  minutes.  Again  allow  dilatometer  to 
cool,  taking  readings  of  temperature  and  volume. 

Does  suspended  transformation  cause  trouble? 

How  does  the  inversion  point  of  sodium  thiosulphate  differ  from 
the  inversion  point  with  Glauber's  salt? 


SUB-GROUP  2 

EUTECTIC  POINTS 

EXPERIMENT  1 
Cryohydric  Points. 

In  this  case  the  problem  is  to  determine  the  conditions  under  which 
solid  solvent  (ice),  solid  solute  (K2SO4),  solution  and  vapor  may 
co-exist.  Under  the  conditions  of  the  experiment,  using  vessels  open 
to  the  air,  the  system  may  not  really  be  in  equilibrium  with  the  vapor 
and  may  be  under  a  pressure  different  from  that  of  the  invariant 
system,  ice,  salt,  solution  and  vapor.  Actually,  however,  the  slight 
and  slowly  acting  readjustments  due  to  these  causes  do  not  have 
much  influence  upon  the  temperature  at  which  ice,  salt  and  solution 
are  in  equilibrium;  and  the  eutectic  temperature  of  a  system  com- 
posed of  non- volatile  or  slightly  volatile  salt,  ice,  solution  and  vapor, 
determined  at  atmospheric  pressure  in  open  vessels,  does  not  differ 
appreciably  from  the  temperature  of  the  system,  salt,  ice,  solution 
and  vapor  in  complete  equilibrium. 

Part  1.  Prepare  a  saturated  solution  of  K2SO4  in  water  and  place 
this  solution  in  a  test  tube  immersed  in  an  ice-salt  freezing  mixture. 
Note  the  temperature  at  one  minute  intervals,  immersing  the 
thermometer  in  the  solution.  Draw  the  usual  curve  between  time 
and  temperature. 

Part  2.  Prepare  a  dilute  solution  of  K2SO4  and  repeat  the  pro- 
cedure of  Part  1  using  5  g.  K2SO4  in  93  cc.  water. 

The  concentration  of  the  solution  at  the  cryohydric  temperature 
may  be  ascertained  by  removing  a  sample  with  a  pipette,  being  care- 
ful to  prevent  the  introduction  of  any  solid  material  into  the  pipette. 
This  sample  may  be  analyzed  and  its  sulphate  content  determined  by 
precipitating  with  barium  chloride. 

EXPERIMENT  2 
Eutectic  Points  by  Cooling  Curves 

By  the  thermometric  method  determine  the  eutectic  point  of  one 
of  the  following  pairs:  naphthalene-anthracene,  naphthalene- 
phenol,  naphthalene-diphenylamine. 

Compare  with  data  and  phase  diagrams  in  LBR. 

63 


SUB-GROUP  3 

TWO  LIQUID  LAYERS 

EXPERIMENT  1 

Melting  under  the  Solvent.  Add  an  excess  of  para-toluidine  to 
water  in  a  test  tube.  Heat  on  a  steam  bath  to  45°  C.  and  note  what 
happens.  At  what  temperature  does  the  para-toluidine  melt? 
What  is  the  melting  point  of  pure  paratoluidine?  FPR,  129  (1917). 

EXPERIMENT  2 

Phenol  and  Water.  Make  up  mixtures  of  phenol  and  water  of  the 
following  composition  in  parts  of  phenol  in  100  parts  of  mixture: 
5,  8,  10,  20,  30,  40,  50,  60,  70,  80,  90. 

Weigh  the  required  amounts  of  phenol  out  as  quickly  as  possible 
to  prevent  absorption  of  moisture  from  the  air.  Let  the  combined 
weight  of  phenol  and  water  in  each  mixture  be  15  or  20  g.  Add  the 
required  amount  of  water  from  a  burette  and  immediately  close  the 
mouth  of  the  test  tube  with  a  cork. 

Beginning  with  the  mixture  containing  10  per  cent  of  phenol,  heat 
each  succeeding  mixture  (up  to  the  90  per  cent  one)  by  immersing  the 
test  tube  in  a  water  bath  (i.  e.  a  beaker).  Place  a  thermometer  in 
the  test  tube  and  stir  thoroughly.  Stirring  by  means  of  a  slow  stream 
of  air  is  very  effective.  When  the  two  layers  disappear,  and  the 
liquid  becomes  homogeneous,  observe  the  temperature. 

Next  remove  the  test  tube  from  the  bath,  and  with  constant  stirring 
and  slow  cooling,  observe  the  temperature  at  which  the  two  layers 
reappear,  i.  e.  when  the  solution  becomes  milky. 

Next  place  the  test-tube  in  a  freezing  mixture  and  determine  the 
temperature  at  which  the  phenol  solidifies  under  the  solution.  Is  this 
the  same  temperature  as  the  eutectic  point?  Explain.  Determine 
the  eutectic  point. 

Draw  a  curve  with  concentrations  as  abscissae  and  temperatures  as 
ordinates. 

The  8  (and  perhaps  the  70)  per  cent  solutions  should  be  homogen- 
eous at  ordinary  temperatures.  On  immersion  in  cold  water,  how- 
ever, the  liquid  layers  will  be  formed  just  as  in  the  other  cases. 
Determine  at  what  temperature  this  occurs. 

The  5,  80  and  90  per  cent  mixtures  should  also  be  homogeneous  at 
room  temperature.  On  cooling  in  a  freezing  mixture,  these  solutions 
do  not  separate  into  two  liquid  layers  but  deposit  a  solid  phase. 
Determine  the  temperature  at  which  solid  first  begins  to  appear  and 
ascertain  the  nature  of  the  solid  phase.  Ice  or  phenol? 

Note.  Do  not  throw  away  the  phenol-water  mixtures  but  return 
them  to  the  bottle  marked  "phenol  residues." 

Note  the  following: 

Take  a  30  per  cent  mixture  of  phenol  in  water  and  heat  to  about 
75°  C.  At  this  temperature,  add  5  to  10  grams  of  solid  phenol. 
Do  two  liquid  layers  form?  Allow  the  solution  to  cool  down  until 
the  layers  appear,  noting  the  temperature.  Represent  what  you  did 
graphically  on  the  curve  obtained  in  Experiment  2. 

64 


Precaution.  Phenol  is  very  corrosive.  Do  not  let  it  remain  in 
contact  with  the  skin. 

References.  LBR,  592;  Lehfeldt,  228;  Rothmund:  Die 
Loslichkeit  (1907). 

EXPERIMENT  3 
Sulphur  and  Aniline.     (Optional) 

Proceeding  exactly  as  you  did  in  the  case  of  phenol  and  water  make 
up  the  following  mixtures  of  sulphur  and  aniline: 

25,  40,  50,  60,  70,  80,  85,  90,  93  per  cent  sulphur. 

Determine  the  temperatures  at  which  the  layers  appear  (i.  e.  the 
clear  liquid  becomes  turbid)  on  cooling  the  clear  solutions  from  a 
temperature  of  140-160°  C.  The  turbidity  will  be  noticed  between 
the  temperature  limits  of  102°  and  140°  C.  Stir  vigorously.  Also 
ascertain  at  what  temperature  the  pure  sulphur  melts  and  at  what 
temperature  it  melts  under  the  solvent.  To  do  this  note  the  tem- 
perature at  which  the  lower  layer  of  aniline  in  sulphur  in  one  of  the 
above  mixtures  solidifies  to  a  crystalline  yellow  mass. 

Draw  a  curve  between  temperature  and  composition. 

Precaution.     Use  roll  sulphur  (not  flowers  of  sulphur). 
Reference.     LBR,  595. 

EXPERIMENT  4 
Three  Components — Chloroform,  Acetic  Acid,  and  Water 

Make  up  mixtures  of  chloroform  and  water  of  the  following  compo- 
sition (by  weight) :  98,  95,  90,  80,  70,  60,  50,  40,  30,  20,  10,  5,  2  parts 
of  chloroform  in  100  of  mixture.  Total  weight  of  each  mixture  to 
be  40  grams. 

Mix  in  100  cc.  glass  stoppered  bottles,  shake  vigorously,  heat  to 
about  40°  in  a  water  bath,  cool  and  allow  to  come  to  equilibrium  by 
standing  a  week. 

When  this  has  been  done  and  the  bottles  are  at  the  same  tempera- 
ture (record)  add  glacial  acetic  acid  from  a  burette  until  a  homo- 
geneous (non-cloudy)  solution  is  obtained.  Shake  constantly  during 
the  addition  of  the  acid.  Calculate  the  weight  of  acetic  acid  neces- 
sary to  produce  a  homogeneous  solution  and  plot  your  results  upon  a 
triangular  diagram. 

Reference.     FPR,  249  (1917). 


SUB-GROUP  4 

PREPARATION  OF  COMPOUNDS 

The  object  of  this  set  of  experiments  is  to  give  practice  in  applying 
phase  rule  methods  to  the  preparation  of  compounds  by  the  system- 
atic use  of  temperature-composition  diagrams. 

65 


EXPERIMENT  1 
Hexahydrate  of  Calcium  Chloride 

Diagram  in  FPR,  155  (1917).  Plan  your  procedure  carefully  and 
report  it  to  the  Instructor  before  doing  this  experiment.  Follow  this 
plan  throughout. 

Note.  Filter  the  CaCl2  solution,  as  it  may  be  turbid  on  account  of 
basic  chlorides  if  made  from  desiccated  CaCl2.  Show  the  compound 
to  the  Instructor. 

EXPERIMENT  2 
Hydrates  of  Potassium  Hydroxide 
Part  1.     Prepare  the  monohydrate  of  KOH. 
Part  2.     Prepare  the  dihydrate  of  KOH. 
Show  the  crystals  to  the  Instructor. 

References.  Pickering:  Jour.  Chem.  Soc.,  63,  899  (1893). 
Note  properties  of  crystals,  898.  Complete  data  and  diagram  in 
LBR,  477. 

EXPERIMENT  3 
Monohydrate  of  Sulphuric  Acid 

Prepare  the  monohydrate  of  H2SO4. 

References.  Pickering:  Jour.  Chem.  Soc.,  57,  338  (1890). 
Complete  data  (SO3  and  water)  and  diagram  in  LBR,  493. 

Hint.  Since  the  solubility  curve  for  the  compound  H2SO4'  H2O 
passes  through  a  very  sharp  maximum  in  respect  to  temperature, 
unless  the  concentration  of  the  solution  is  very  accurately  adjusted 
to  be  equal  to  that  of  the  maximum  point,  one  is  very  apt  to  meet 
with  failure  unless  the  solution  is  cooled  to  a  very  low  temperature. 
Prepare  the  solution  and  divide  it  into  two  equal  parts.  _  Try  to 
crystallize  out  the  monohydrate.  If  you  fail,  the  solution  is  either 
too  concentrated  or  too  dilute  (unless  supersaturation  has  caused  the 
trouble).  To  one  of  the  tubes  add  a  drop  of  water,  to  the  other  a 
drop  of  concentrated  acid  and  again  attempt  to  crystallize  the  mono- 
hydrate.  Continue  this  procedure  until  you  succeed.  Show  the 
crystals  to  the  Instructor  and  record  the  temperature  at .  which  the 
last  crystals  disappear  on  warming. 

EXPERIMENT  4 
Carnallite— KC1  -MgCl2  '6H2O 
Discussion. 

Cf .  FPR,  280-298  (1917) ;  see  isothermal  diagram  for  25°  C.  in 
Whetham:  Solutions,  404  (1902);  excellent  discussion  by  Hilde- 
brand:  Jour.  Ind.  Eng.  Chem.,  10,  97  (1918). 

It  is  obvious  that  if  one  prepares  a  solution  containing  equimole- 
cular  quantities  of  KC1  and  MgCl2'6H2O  and  evaporates  until  the 

66 


solution  phase  just  disappears,  carnallite  will  be  formed,  since  this 
salt  is  stable  above  — 21°.  This  method  however  is  not  elegant  and 
if  the  evaporation  is  discontinued  at  any  point  short  of  complete 
disappearance  of  the  liquid  phase  a  mixture  of  carnallite  and  KC1  will 
be  obtained.  It  is  important  to  remember  that  carnallite  cannot  be 
in  equilibrium  with  a  solution  containing  MgCl2  and  KC1  in  the  ratio 
of  1:1.  When  carnallite  is  dissolved  in  water  the  solution  soon 
becomes  saturated  with  KC1  and  this  salt  is  precipitated  while 
carnallite  continues  to  dissolve.  It  is  not  until  the  MgCl2  content 
of  the  solution  rises  to  a  high  value  by  the  precipitation  of  KC1,  that 
carnallite  can  exist  as  stable  phase  in  contact  with  solution. 

Procedure. 

Prepare  a  solution  of  MgCl2  and  KC1  in  the  proper  molecular  ratio 
to  insure  the  separation  of  carnallite  as  the  first  solid  phase 
on  cooling  or  dehydrating.  Show  the  crystals  to  the  Instructor. 
Prove  that  they  really  are  carnallite. 

For  data  regarding  the  composition  of  the  solution  cf.  FPR,  298, 
296  (1917). 

Suggest  a  simple  method  of  obtaining  KC1  from  Stassfurt  carnallite. 

EXPERIMENT  5 
Copper-potassium  Chloride 

Following  the  procedure  used  in  preparing  carnallite,  make  the 
blue  double  salt.  Test  for  purity  by  determining  the  inversion  point 
for  the  breakdown,  2  KC1  •  CuCl2  •  2H2O  -+  KC1 ;  CuCl2  +  KC1  +  2H2O 
If  the  salt  is  green  the  result  is  not  entirely  satisfactory. 

Note.  2  KCl-CuCl2-2H2O,  like  carnallite,  is  unstable  in  contact 
with  solution  containing  KC1  and  CuCl2  in  the  ratio  2:1,  but  is  stable 
in  contact  with  a  solution  containing  these  salts  in  the  ratio  1 :1  or  1 :2. 
Bancroft:  Phase  Rule,  176  (1897). 


EXPERIMENT  6  (OPTIONAL) 
Lead  Potassium  Iodide 

Prepare  lead  potassium  iodide,  PbI2'KI'2H2O.  Bancroft,  179 
(1897);  Abegg:  Handbuch,  III  (2)  667;  Schreinemakers:  Zeit.  phys. 
Chem.  10,  467  (1892). 

Note.  Schreinemakers'  diagram  on  page  471  indicates  that  the 
double  salt  is  stable  only  in  a  solution  containing  KC1  in  excess. 

EXPERIMENT  7  (OPTIONAL) 
Astracanite,  Na2SO4-  MgSO4-  4H2O 

Reference.     FPR,  264  (1911). 

Report  results.     Write  the  reaction. 

67 


SUB-GROUP  5 

INDIRECT  ANALYSIS 
Discussion. 

Under  some  circumstances  solid  separates  out  from  a  liquid  phase 
in  a  form  which  renders  direct  analysis  very  difficult  and  uncertain. 
The  solid  may  be  unstable  and  it  may  be  impossible  to  remove  adher- 
ing mother-liquor.  Indirect  analysis  is  then  resorted  to.  Many 
methods  of  indirect  analysis  have  been  proposed;  the  following 
experiment  illustrates  one  of  the  most  satisfactory. 

References.     Bancroft:     Jour.  Phys.  Chem.,  6,  178  (1902). 
Browne:     Ibid,  6,  281  (1902). 
FPR,  236,  310  (1917). 


EXPERIMENT  1 

Determination  of  Solid  Phases 
Discussion. 

Let  us  suppose  a  system  to  be  composed  of  three  components  A,  B, 
and  C,  all  of  them  miscible  in  the  liquid  phase.  Starting  with  a 
system  composed  of  the  homogeneous  (unsaturated)  solution  in  con- 
tact with  vapor,  let  the  composition  of  the  solution  be  a  per  cent 
of  A,  b  per  cent  of  B,  and  c  percent  of  C. 

Next,  without  changing  the  total  amount  of  A,  B,  and  C  in  the 
system  (no  loss  by  evaporation,  etc.)  cool  until  a  single  solid  phase 
separates  out  and  the  system  solid-liquid  is  produced.  Suppose  that 
a  qualitative  analysis  of  the  solid  phase  indicates  that  C  is  not  present 
in  the  solid.  There  are  three  possibilities,  as  follows: 

(1)  Solid  is  pure  A  or  pure  B. 

(2)  Solid  is  a  compound  of  A  and  B. 

(3)  Solid  is  a  solid  solution  of  A  and  B  or  an  absorption  compound. 
Without,  removing  the  solid,  pipette  out  some  of  the  clear  mother- 
liquor  and  analyze  it.     Let  the  composition  now  be  (in  per  cent)  a', 
b',  and  c'.     The  following  relations  hold  true  for  the  two  solutions: 

a  +b  +c   =  100  (1) 

a'  -f  b'  +  c'  =  100  (2) 

Next  divide  (2)  by  ^-,  whence 


vSince  C  has  not  separated  out  in  the  solid  phase  and  the  total 
amount  of  C  in  the  liquid  phase  therefore  remains  unchanged,  the 
composition  of  the  solid  phase  must  be  proportional  to  (FPR,  232)  : 


68 


If  MA  and  M_,  are  the  respective  molecular  weights,  then  the 

A.  13 

molecular  composition  of  the  solid  phase  is  given  by  the  expression 
ao'-a'A  /bc-jyc 


C'MA  ;       ^  ,  MB 

From  (5)  the  number  of  molecules  of  B  combined  with  one  mole- 
cule of  A  becomes 


_^A      (bc>  -  b'A 
M        ^ac'  -  a'c/ 


B 
Procedure. 

Prepare  a  solution  of  50  g;  sodium  sulphate  decahydrate  (Glauber's 
salt)  and  10  g.  sodium  chloride  in  100  cc.  of  distilled  water.  Filter 
the  hot  solution.  Cool  to  45°  C.  and  analyze  the  solution  for  sodium 
chloride  and  sodium  sulphate.  See  below  for  procedure.  Keep  the 
solution  in  a  stoppered  flask  or  Erlenmeyer.  Run  in  duplicate. 

Cool  the  solution  until  solid  crystallizes  out  in  considerable  amount, 
then  carefully  pipette  two  samples  of  the  solution  for  analysis.  It 
may  be  found  advisable  to  fit  to  the  end  of  the  pipette  a  bit  of  glass 
tubing  containing  glass  wool  or  cotton  to  serve  as  a  filter.  Separate 
some  of  the  solid  and  wash  with  a  very  little  water.  Has  any  sodium 
chloride  been  precipitated? 

Analysis.  Determine  NaCl  in  one  sample  (1  g.)  with  standard 
silver  nitrate  (shelf)  using  K2CrO4  as  indicator.  Evaporate  a 
second  sample  to  dryness  (being  careful  to  avoid  spattering)  and 
determine  total  chloride  and  sulphate.  Determine  water  by 
difference.  Using  equation  (6)  determine  the  chemical  formula  of 
the  solid  phase,  assuming  that  no  solid  solutions  are  formed  in  this 
experiment. 

How  else  might  one  determine  approximately  the  composition  of 
the  solid  in  the  above  experiment,  using,  of  course,  an  indirect 
method? 

Outline  the  procedure  in  case  component  C  also  separates  out  in 
the  solid  phase.  See  references  (Triangular  Diagrams). 

How  could  one  distinguish  between  compound  and  solid  solution? 

Why  must  the  two  salts  have  an  ion  in  common? 

How  can  one  tell  whether  the  number  of  solid  phases  precipitated 
from  the  solution  is  one  or  two? 


EXPERIMENTAL  GROUP  XVI 

COLLOID  CHEMISTRY 

This  comprehensive  group  of  experiments  serves  to  illustrate  some 
of  the  more  important  and  interesting  properties  of  colloidal  systems. 
Typical  colloids  are  prepared  and  studied,  particularly  from  the 
point  of  view  of  Bancroft:  Jour.  Phys.  Chem.,  18,  549  (1914). 
Read  the  article  before  beginning  experimental  work  in  this  group. 

General  Texts  in  Colloid  Chemistry. 

Alexander:     Colloid  Chemistry  (1919). 
Bancroft:     Applied  Colloid  Chemistry  (1920). 
Burton:     Physical  Properties  of  Colloidal  Solutions  (1916). 
Cassuto:     Der  Kolloide  Zustand  der  Materie  (1911). 
Freundlich:     Kapillarchemie  (1909). 

Hatschek:     An  Introd.  to  the  Physics  and  Chemistry  of  Colloids 
(1919). 
Miiller:     Chemie  der  Kolloide  (1907). 

Ostwald  (w°) :     Grundriss  der  Kolloidchemie  (1911-12). 

Ostwald  (w°)  (Fischer):  Theoretical  and  Applied  Colloid 
Chemistry  (1915). 

Ostwald  (w°)  (Fischer) :   Handbook  of  Colloid  Chemistry  (1915) 

Svedberg:  Die  Methoden  zur  Herstellung  kolloider  Losungen 
usw.  (1909). 

Taylor:     The  Chemistry  of  Colloids  (1915). 

Willows  and  Hatschek:     Surface  Tension  (1915). 

Zsigmondy  (Alexander) :     Colloids  and  the  Ultramicroscope  (1909) . 

Zsigmondy:     Kolloidchemie  (1912). 

Zsigmondy  (Spear):     Colloid  chemistry  (1917). 

Journals 

Journal  of  Physical  Chemistry,     (special  articles). 
Kolloidchemische  Beihefte  (special  articles). 
Kolloid-Zeitschrift.     (1906—). 

SUB-GROUP     1 

DIFFUSION,  DIALYSIS  AND  MEMBRANES 

EXPERIMENT  1 
Diffusion  of  Solutions. 

Obtain  six  test  tubes,  fitting  each  with  a  rubber  stopper  (one  hole), 
and  prepare  six  15  cm.  lengths  of  narrow-bore  (2.  5-3  mm.  internal 
diam.)  glass  tubing.  Seal  one  end  of  each  length  of  tubing  and  fill 

70 


completely  with  distilled  water.  Place  10  cc.  of  solution  to  be  tested 
in  each  test  tube,  insert  a  water-filled  diffusion  tube  in  the  stopper 
and  place  it  in  the  test  tube,  immersing  open  end  of  the  diffusion 
tube  just  below  the  surface  of  the  solution.  Work  carefully.  Set  aside 
the  test  tubes  in  a  safe  place  and  make  observations  at  regular  inter- 
vals, recording  the  time.  Test  the  following  solutions: 

KMnO4  solution  —  N/50. 

KMnO4  solution  —  N/5. 

Congo  red  — 1/5  of  one  per  cent. 

Methyl  violet  or  safranine  — 1/5  of  one  per  cent. 

Arsenious  sulphide  sol. — (See  Part  3  below). 

Ferric  oxide  sol. — (Sse  Part  3  below). 

Optional  Method.  The  following  experiments  are  similar  to  those 
of  Graham.  A  small,  two-dram  vial  is  fastened  to  the  bottom  of  a 
tall,  narrow  beaker  (250  cc.  capacity)  by  means  of  paraffin. 

Fill  the  vial  carefully  with  the  solution  containing  the  solute  whose 
rate  of  diffusion  is  to  be  measured  and  cover  it  securely  with  a  small 
cover-glass  (20  millimeters).  Be  sure  that  no  solution  is  spilled 
from  the  vial  during  the  process  of  filling  and  covering.  Pour  dis- 
tilled water  into  the  beaker  until  it  is  nearly  full  and  the  vial  is  well 
covered,  taking  care  to  have  the  water  level  at  the  same  height  in 
each  beaker.  Finally,  slide  the  cover  glass  carefully  off  the  mouth  of 
the  vial  by  means  of  a  clean  glass  rod. 

A  two  cc.  test-sample  is  then  pipetted  from  the  liquid  in  the  beaker 
at  a  point  about  three  centimeters  above  the  open  mouth  of  the  vial. 
Mark  this  position  by  means  of  a  label  placed  on  the  wall  of  the 
beaker.  Be  careful  not  to  stir  the  liquid.  Test  for  chlorine  as  ion 
with  silver  nitrate  making  a  rough  nephelometric  estimation  of  the 
relative  amounts  of  silver  chloride  formed  in  each  sample.  Test  for 
organic  matter  by  evaporating  a  test  sample  to  dryness  in  a  clean 
porcelain  dish  and  carbonizing  the  residue. 

It  is  essential  that  the  water  levels  be  the  same  in  each  beaker,  that 
the  sample  be  pipetted  from  equal  distances  above  the  mouth  of  the 
vial  and  that  the  beakers  and  solution  remain  absolutely  undisturbed. 
Withdraw  test  samples  at  the  beginning  and  after  1,  2,  4  and  7  days, 
noting  the  exact  time. 

The  following  solutions  are  to  be  tested: 

(1)  One  per  cent  solution  of  gelatine. 

(2)  Five  per  cent  solution  of  sodium  chloride. 

(3)  Twenty-five  per  cent  solution  of  sodium  chloride. 

Note.  Prepare  a  5  per  cent  solution  of  gelatine  for  this  and  subse- 
quent work  as  follows:  Soak  2  g.  of  gelatine  in  cold  water  until  soft, 
pour  off. the  water  and  to  the  softened  gelatine  add  enough  warm 
water  to  make  about  40  cc.  of  solution.  On  cooling,  a  jelly  will  form 
which  readily  melts  when  the  beaker  with  the  jelly  is  warmed  on  the 
steam  bath.  Do  not  warm  over  a  flame  as  the  beaker  will  almost 
certainly  crack.  Dilute  the  gelatine  solution  as  required. 

71 


EXPERIMENT  2 
Diffusion  Through  a  Jelly 

Obtain  eight  small  test  tubes  and  fill  each  half  full  of  liquid  5  per 
cent  gelatine  and  allow  this  to  solidify.  Pour  into  the  tubes,  on  top 
of  the  gelatine,  the  solutions  or  sols  specified  below,  being  careful 
that  the  latter  are  cold  so  that  they  do  not  liquefy  the  jelly. 

If  they  diffuse,  the  substances  in  solution  will  tend  to  pass  from  the 
upper  aqueous  layer  into  the  lower  portion  occupied  by  the  gelatine 
and  the  process  may  be  observed  by  means  of  the  coloration  produced 
in  the  jelly.  If  the  colored  substance  forms  a  true  solution,  the 
diffusion  of  the  solute  through  a  jelly  occurs  almost  as  rapidly  as 
through  pure  water  itself.  On  the  other  hand,  colloidal  solutions 
show  practically  no  evidence  of  diffusion.  We  may,  therefore,  dis- 
tinguish between  the  two  classes  of  solution  by  means  of  this  method, 
provided  the  jelly  is  not  "semi-permeable"  to  the  dissolved  solute. 

Observe  the  condition  of  each  tube  after  twenty-four  hours  and 
again  after  a  week.  Keep  the  tubes  in  a  cool  place.  Use  the  follow- 
ing solutions  (shelf) : 


(1)  Eosine 
(2)  Congo  red 
(3)  Safranine 
(4)  Picric  acid 
(5)  Methylene  blue 
(6)  Arsenious  sulphide  sol 

1/5  of  one  per  cent. 
1/5  of  one  per  cent. 
1  /5  of  one  per  cent. 
1/5  of  one  per  cent. 
1/5  of  one  per  cent, 
(see  below). 

-(7)  Ferric 'oxide  sol  (see- below). 
(8)  Mixture  Congo  red  and  picric  acid,  picric  acid  in  excess. 

From  the  data  obtained  in  these  experiments  what  conclusion  do 
you  draw  regarding  the  nature  of  the  above  solutions? 

EXPERIMENT  3 

Dialysis  v/ith  Collodion.  Instead  of  using  parchment,  prepare 
collodion  dialyzing  tubes  as  follows:  Take  one  of  the  inner  test 
tubes  of  heavy  glass  used  in  the  free2ing  point  determinations  and 
wet  the  inner  walls  completely  with  a  fairly  thick  film  of  collodion 
solution  (soluble  cotton  in  a  mixture  of  ether  and  alcohol) .  Do  this 
quickly  while  spinning  the  tube  to  make  the  collodion  film  uniform. 

As  soon  as  the  collodion  "sets"  blow  air  into  the  tube  to  remove  the 
ether.  This  process  should  take  about  five  minutes.  Then  pour 
water  into  the  test  tube  and  gradually  loosen  the  collodion  from  the 
glass.  With  moderately  careful  manipulation,  a  transparent,  tough 
dialyzing  tube  can  be  obtained  which  is  more  convenient  and  less 
expensive  than  the  parchment  dialyzers  ordinarily  used.  Having 
prepared  the  tube,  test  for  leaks  by  filling  with  water  and  if  intact, 
immerse  completely  in  a  large  beaker  of  water  to  remove  the  alcohol. 
Soak  until  the  next  period,  changing  the  water  from  tirn.e  to  time. 
Make  three  dialyzing  tubes. 

Fill  one  nearly  full  with  a  mixsd  solution  containing  1  per  cent 
gelatine  plus  5  per  cent  of  sodium  chloride.  Place  this  in  a  beaker  of 
distilled  water  and  test  the  water  at  stated  interval  for  NaCl  and 
.gelatine. 

72 


Fill  the  second  tube  with  a  solution  of  safranine.  Place  this  in  a 
second  beaker  of  water  and  observe  diffusion.  In  the  third  tube  place 
a  solution  of  Congo  red.  Does  this  diffuse? 

EXPERIMENT  4 
Semipermeable  Membranes 

Into  a  small  bottle  pour,  very  carefully  and  in  the  order  given,  the 
following  liquids:  Chloroform,  water,  and  ether.  Three  layers 
should  be  present.  Note  the  thickness  in  mm.  of  each  layer. 

Let  the  bottle  stand  undisturbed  for  a  week  and  again  measure 
the  thickness  of  the  layers.  Continue  the  experiment  until  one  of  the 
three  original  layers  disappears.  Explain. 

Reference.     Kahlenberg:     Jour.  Phys.  Chem.,  10,  146  (1906). 

EXPERIMENT  5 
Osmosis  and  Semipermeable  Membranes 

Part  1.  Fill  a  test  tube  with  a  M/2  CuSO4,  then,  by  means  of  a 
pipette  placed  in  this  solution  add  slowly  and  carefully  a  small 
amount  of  M/2  potassium  ferrocyanide.  A  globule  should  form, 
consisting  of  the  solution  of  ferrocyanide  surrounded  by  a  gelatinous 
membrane  of  brown  copper  ferrocyanide.  Carefully  detach  the 
globule  from  the  end  of  the  pipette  and  it  will  sink,  owing  to  the 
greater  density  of  the  ferrocyanide  solution. 

Observe  carefully  any  changes  that  may  occur  in  the  copper 
sulphate  solution  surrounding  the  globule.  Set  aside  the  test  tube 
and  keep  it  constantly  under  observation.  What  happens? 
Explain. 

Part  2.  Plant-like  Growths.  Fill  a  small  beaker  with  dilute 
sodium  silicate  (water  glass)  solution  and  drop  into  the  liquid  one  or 
two  crystals  each  of  CuSO4,  MnSO4,  CoSO4,  etc.  What  happens? 
Explain. 

SUB-GROUP  2 

ADSORPTION 

The  following  experiments  are  designed  to  illustrate  adsorption 
phenomena.  Adsorption  is  the  basis  of  colloid  chemistry.  All  the 
experiments  of  Sub-groups  3  and  4  illustrate  this  point. 

EXPERIMENT  1 
Adsorption  by  Bone  Black 
Part  1.     Boil  a  dilute  solution  of  litmus  with  bone  black.     Filter. 

Part  2.  Repeat,  using  dilute  solution  of  indigo.  Are  the  colors 
removed?  Explain. 

Part  3.  Prepare  a  dilute  solution  of  silver  nitrate.  Divide  this 
into  two  portions.  To  one  portion  add  about  one-tenth  its  volume  of 
bone  black  and  shake  vigorously  for  at  least  three  minutes.  Then 

73 


filter  and  add  NaCl  to  both  portions.      Compare  the  amounts  of 
precipitated  silver  chloride. 

Bone  black  or  animal  charcoal  contains  85  per  cent. of  calcium 
phosphate  and  about  15  per  cent  of  carbon. 

EXPERIMENT  2 
Selective  Adsorption 

Part  1.  Prepare  about  250  cc.  of  indicator  solution  as  follows: 
To  250  cc.  of  distilled  water  add  a  little  phenolphthalein  and  a  trace 
of  NaOH,  just  enough  to  color  the  liquid  pink. 

Part  2.  Ina  test  tube  shake  fuller's  earth  with  distilled  water  and 
add  some  of  this  muddy  suspension  to  one  of  the  test  tubes  containing 
the  indicator.  Is  the  color  removed? 

Part  3.  Allow  this  muddy  suspension  to  settle  and  then  add  the 
supernatant  clear  liquid  to  a  second  test  tube  colored  with  indicator. 
Filter  the  supernatant  liquid  to  remove  all  the  fuller's  earth.  Is  this 
filtered  liquid  acid? 

Part  4.  Moisten  a  little  fuller's  earth  with  boiled  water  and  test 
with  blue  litmus  by  pressing  the  latter  down  on  the  earth. 

Reference.     Cameron:     Jour.  Phys.  Chem.,  14,  400  (1910). 

Part  5.  Add  blue  litmus  solution  to  fuller's  earth  suspended  in 
water.  Notice  the  change. 

Part  6.  Add  some  fuller's  earth  to  a  dilute  solution  of  methyl 
violet  and  shake.  Filter,  noting  color  of  filtrate  and  of  earth.  Is 
the  color  removed  from  the  earth  by  water  or  alcohol? 

Part  7.  Repeat  the  last  experiment,  using  eosin  instead  of  methyl 
violet.  Note  any  differences  in  behavior. 

Part  8.  Moisten  some  absorbent  cotton  with  freshly  boiled  water 
(free  from  CO2)  and  wrap  it  around  a  strip  of  blue  litmus  paper. 
For  comparison  of  the  original  and  the  final  color,  let  about  half  an 
inch  of  the  paper  protrude  beyond  the  cotton.  Explain  your  results. 
Compare  Part  4,  above. 

EXPERIMENT  3 
Adsorption  by  Iron  Oxide.     The  Antidote  for  Arsenic  Poisoning 

Hydrous  ferric  oxide  is  precipitated  from  a  solution  of  ferric 
sulphate  or  chloride  by  adding  an  excess  of  magnesia.  Shake 
vigorously.  Then  prepare  a  dilute  solution  of  As2O3  and  filter,  and 
test  the  filtrate  for  arsenic  with  H2S. 

Be  sure  that  the  As2O3  solution  is  very  dilute.  Test  half  the 
original  solution  with  H2S  for  arsenic.  Only  a  slight  test  should  be 
obtained,  if  the  experiment  is  to  work  well.  Then  test  the  second 
half  of  the  As2O3  solution  after  treatment  with  the  ferric  hydroxide 
mixture.  Has  the  arsenic  been  adsorbed?  Should  the  arsenic  be 
completely  adsorbed?  Explain. 

74 


EXPERIMENT  4 
Adsorption  Compounds.     Carey  Lea's  "Photohalides" 

Reference.     Carey  Lea:     Am.  Jour.  Science,   (3)  34,  349,  480, 

489  (1887). 

Method  suggested  by  Luppo-Cramer :  Kolloid-Zeit.,  2,  360 
(1908). 

To  3.5  cc.  of  10  per  cent  KBr  add  5.5  cc.  of  10  per  cent  AgNO3. 
To  this  mixture  containing  AgBr  plus  AgNO3  in  excess  add  the 
following  solution: 

7.5  cc.  Rochelle  salts  (1:3)  plus  2.5  cc.  of  ferrous  sulphate  (1:3). 
Do  not  add  the  Rochelle  salts  and  ferrous  sulphate  solutions  sepa- 
rately. 

Wash  the  dark  colored  precipitate  several  times  by  decantation  and 
finally  with  a  mixture  of  equal  parts  concentrated  HNO3  (1.4  sp.gr.) 
and  water.  An  intense  blue-violet  color  should  develop. 

The  photohalides  of  silver  are  adsorption  compounds  of  silver  with 
silver  chloride  and  are  similar  to  the  "subsalts"  of  silver  composing 
the  "latent  image"  in  an  exposed  photographic  plate. 

EXPERIMENT  5 
Selective  Adsorption  and  Capillary  Diffusion 

Part  1.  Place  several  drops  of  a  mixed  solution  of  CuSO4  and 
CdSO4  (shelf)  on  the  center  of  a  square  of  blotting  paper  (6  by  6  in.). 
Allow  the  drops  to  diffuse  until  a  large  round  spot  has  formed,  then 
hold  the  paper  in  a  stream  of  H2S  gas.  Which  "diffuses"  farthest, 
water,  CuSO4,  or  CdSO4?  Cf.  Gordon:  Jour.  Phys.  Chem.,  18, 
337  (1914). 

Part  2.  Suspend  strips  of  blotting  paper  (1  cm.  broad  and  20  cm. 
long)  in  water  solutions  of  the  following  substances:  Congo  red; 
picric  acid;  cosin;  methylene  blue;  methylene  blue  plus  cosin. 
Note  the  height  to  which  the  water  and  dye  rise. 

Reference.     Goppelsroeder:     Kapillaranalyse  (1906). 

EXPERIMENT  6 
Adsorbed  Air  in  Charcoal 

Fit  a  cylinder  (100  cc.)  with  a  three-hole  rubber  stopper.  Into  one 
hole  introduce  the  delivery  tube  of  a  burette  filled  with  water.  In 
the  second  place  a  thermometer.  In  the  third  place  a  glass  tube  lead- 
ing to  a  water- filled  graduated  cylinder  (capacity  250  cc.)  inverted 
over  water  in  a  trough. 

Place  a  volume  of  50  apparent  cc.  of  granular  cocoanut  charcoal  in 
the  cylinder.  Then  add  water  slowly  from  the  burette,  recording  the 
volume  added.  Continue  to  add  water  until  its  level  rises  to  the 
surface  of  the  charcoal.  Measure  the  volume  of  air  displaced. 

Take  the  temperature  before  and  after  adding  the  water.  Have 
the  water  in  the  burette  and  the  charcoal  at  the  same  temperature  in 
the  beginning. 

75 


SUB-GROUP  3 

PEPTIZATION 

EXPERIMENT  1 

Peptization  by  Adsorbed  Ions 

Lottermoser:  Jour.  Praktische  Chemie,  [2]  72,  39  (1905);  73, 
374(1906);  Zsigmpndy  (Spear)  179;  Ostwald  (Fischer) :  Theoretical 
and  Applied  Colloidchemistry,  115. 

Part  1.  Prepare  a  small  quantity  of  silver  bromide  and  wash  the 
precipitate  thoroughly  by  decantation.  Place  approximately  equal 
amounts  of  the  freshly  prepared  silver  bromide  in  each  of  five 
stoppered  test  tubes.  In  the  first  test  tube  place  distilled  water  (10 
cc.);  in  the  second,  N/100  KBr;  in  the  third,  N/30  KBr;  in  the 
fourth,  N/10  KBr,  and  in  the  fifth,  N/5  KBr.  Shake  thoroughly  and 
after  allowing  the  test  tubes  to  remain  standing  several  minutes, 
describe  the  appearance  of  each  tube.  In  which  is  the  supernatant 
liquid  most  turbid? 

The  process  constitutes  a  dispersion  method  of  preparing  colloidal 
silver  bromide. 

Part  2.  Fill  two  burettes  with  N/20  AgNO3  and  N/20  NH4CNS 
(shelf).  Fit  a  small  Erlenmeyer  flask  with  a  solid  rubber  st.opper. 

Perform  the  following  experiments: 

(a)  To  10  cc.  AgNO3  in  flask  add  quickly  10  cc.  NH4CNS,  stopper 

and  shake. 

(b)  To  10  cc.  AgNO3  in  flask  add  quickly  12  cc.  NH4CNS,  stopper 

and  shake. 

(c)  To  10  cc.  NH4CNS  in  flask  add  quickly  10  cc.  AgNO3,  stopper 

and  shake. 

(d)  To  10  cc.  NH4CNS  in  flask  add  quickly  12  cc.  AgNO3,  stopper 

and  shake. 

What  striking  differences  do  you  observe  and  how  do  you  account 
for  them? 

(e)  Refill  the  burettes  and,  placing  10  cc.  AgNO3  in  a  flask  run  in 

NH4  CNS  from  a  burette  (not  too  rapidly)  until  floccula- 
tion  occurs.  Shake  and  note  the  volume  of  NH4CNS 
added.  Repeat  adding  NH4CNS  more  slowly  as  end- 
point  is  reached.  The  end-point  represents  the  isoelectric 
point  (define). 

(f)  Place  10  cc.  NH4CNS  in  a  flask  and  add  AgNO3  following 

the  procedure  of  (b)  5  preceding.     Explain. 

EXPERIMENT  2 
Peptization  by  Adsorbed  Colloid 

Prepare  5  per  cent  solutions  of  chromic  and  ferric  chlorides.  Mix 
in  the  proportions  specified  below.  Then  add  10  per  cent  NaOH  in 
excess.  Note  the  color  and  appearance  of  the  precipitate  (if  any) 
and  of  the  supernatant  liquid.  Use  test-tubes  and  shake. 

76 


Ferric  Chloride  Chromic  Chloride  Remarks 

(cc).  (cc.) 

10 0 

8 2 

5 5 

3 7 

2 8  

1 9  

0 10  

Cf.  Nagel:     Jour.  Phys.  Chem.,  19,  331,  569  (1915). 

EXPERIMENT  3 
Peptization  by  Adsorbed  Colloid  (Protective  Colloids) 

Solution  A:     5  cc.  N/2  AgNO3  +  5  cc.  of  5  per  cent  gelatine. 
Solution  B:     5  cc.  N/2  KBr  +  5  cc.  of  5  per  cent  gelatine. 

Part  1.  After  thoroughly  mixing  each  solution,  pour  B  into  A, 
shake  and  note  any  changes.  Place  the  mixture  in  the  sunlight  and 
note  results.  Repeat  the  above  experiment,  replacing  the  gelatine 
solution  by  an  equal  volume  of  pure  water.  Was  AgBr  formed  in  the 
first  experiment  with  gelatine.  How  might  one  prove  this? 

Part  2.  Prepare  some  silver  bromide,  wash  by  decantation  and 
remove  to  a  filter  paper.  Divide  into  two  portions.  Place  one  por- 
tion in  an  air  bath  and  dry  for  an  hour  at  120°,  being  careful  not  to 
exceed  this  temperature. 

To  the  freshly  prepared  moist  silver  bromide  add  a  few  cubic  centi- 
meters of  hot  5  per  cent  gelatine  and  shake  vigorously.  Is  a  suspen- 
sion formed?  Do  the  same  thing  with  the  dried  silver  bromide  and 
note  any  differences  in  its  behavior  compared  with  that  of  the  freshly 
prepared  substance.  What  is  the  effect  of  "ageing?" 

Part  3.  Grind  a  little  roll-sulphur  with  a  5  per  cent  gelatine  solu- 
tion in  a  mortar  until  a  milky  suspension  is  formed.  Pour  some  of 
this  suspension  into  water  and  note  the  color. 


SUB-GROUP  4 
PREPARATION   AND   FLOCCULATION   OF   SUSPENSIONS 

EXPERIMENT  1 
Colloidal  Arsenious  Sulphide  (Condensation  Method) 

In  a  clean  beaker,  boil  about  6  grams  of  As2O3  with  100  cc.  of 
distilled  water  for  fifteen  minutes.  Cool,  filter  and  dilute  to  100  cc. 
Pass  clean  hydrogen  sulphide  gas  into  the  solution  of  arsenious  acid 
until  no  further  action  takes  place.  Remove  excess  of  H2S  by  blow- 
ing a  slow  stream  of  air  through  the  suspension  and  then  filter. 

Describe  the  appearance  of  the  suspension  as  to  color,  turbidity, 
etc.,  and  perform  the  following  tests.  (See  also  diffusion  experi- 
ments) . 

77 


(a)  To  10  cc.  colloidal  As2S3  add  2  cc.  M/20  HC1. 

(b)  To  10  cc.  colloidal  AsaS3  add  2  cc.  M/20  NaCl. 

(c)  To  10  cc.  colloidal  As2S3  add  2  cc.  M/20  MgCl2. 

(d)  To  10  cc.  colloidal  As2S3  add  2  cc.  M/20  Al  (NO3)3. 

(e)  To  10  cc.  colloidal  As2S3  add  2  cc.  M/20  Na2SO4. 

Which  produces  flocculation  most  quickly?     Explain. 

Colloidal  As2S3  thus  prepared  is  a  negative  suspension.  That  is, 
the  particles  of  the  disperse  phase  carry  a  negative  charge  due  to 
preferential  adsorption  of  anions  from  H2S  present  in  solution. 

Place  a  test  tube  containing  10  cc.  of  As2S3  suspension  in  an  ice  salt 
freezing  mixture  until  frozen  solid.  Warm  the  test  tube  gently  until 
the  ice  is  melted.  What  effect  upon  the  suspension  is  noticed? 

EXPERIMENT  2 
Colloidal  Ferric   Oxide    (Condensation  Method) 

Add  about  0.5  gram  of  crystallized  ferric  chloride  to  100  cc.  of 
boiling  distilled  water.  Then  boil  the  solution  gently  for  about  ten 
minutes,  replacing  the  water  boiled  away.  Note  the  color  and  appear- 
ance of  the  hot  solution,  and  compare  with  the  color  of  a  solution 
made  by  adding  FeCl3  to  cold  water.  Explain  the  change.  What 
is  this  process  called? 

(a)  To  10  cc.  of  the  iron  oxide  suspension  add  2  cc.  M/20  NaCl. 

(b)  To  10  cc.  of  the  iron  oxide  suspension  add  2  cc.  M/20  MgCl2. 

(c)  To  10  cc.  of  the  iron  oxide  suspension  add  2  cc.  M/20  Na2SO4. 

(d)  To  10  cc.  of  the  iron  oxide  suspension  add  2  cc.  M/20  citric  acid 

(e)  To  10  cc.  of  the  iron  oxide  suspension  add  trace  of  H2SO4. 

Which  causes  the  most  rapid  flocculation?  Explain.  What  is 
the  precipitate? 

The  ferric  oxide  suspension  as  prepared  above  is  positive. 

Optional  Experiment.    Colloidal  Ferric  Oxide  (Dispersion  Method) 
Reference.     Kratz:     Jour.  Phys.  Chem.,  16,  126  (1912). 
Prepare  Fe2O3  suspension  by  the  method  of  washing  out  the  coagu- 
lating salt,  following  Kratz's  procedure. 

%  EXPERIMENT  3 
Mutual  Flocculation  of  Two  Suspensions 

Study  the  mutual  flocculation  of  colloidal  As2S2  and  Fe2O3/  two 
oppositely  charged  suspensions.  Plan  your  own  experiments. 

EXPERIMENT  4 
Colloidal  Silica.     (Condensation  Method) 

To  10  cc.  of  syrupy  sodium  silicate  solution  add  30  cc.  of  water  and 
pour  the  resulting  solution  into  a  mixture  of  25  cc.  of  concentrated 
hydrochloric  acid  previously  diluted  with  an  equal  volume  of  water. 
A  limpid  mixture  will  result,  consisting  of  a  suspension  of  hydrated 
silica. 

78 


Warm  some  of  this  solution  nearly  to  boiling  and  allow  it  to  stand 
undisturbed  for  a  few  minutes.  What  has  occurred?  Can  the  sus- 
pension be  restored?  Study  the  jelly  obtained.  How  does  it  differ 
from  gelatine  or  agar  agar? 

EXPERIMENT  5 

Colloidal  Metals   (Condensation  Methods) 
Part  1.     Colloidal  Silver.     Gelatine  as  Protecting  Colloid. 

To  5  cc.  of  water  in  a  test  tube  add  about  1  cc.  of  M/10  AgNO3 
solution,  mix  well  and  treat  with  NaOH  in  slight  excess.  What  is 
formed? 

To  5  cc.  of  a  5  per  cent  gelatine  solution  in  a  test  tube  add  about 
1  cc.  M/10  AgNO3,  mix  well  and  treat  with  NaOH  in  slight  excess. 
Note  any  unusual  action.  Then  heat  the  test  tube  until  contents  are 
about  to  boil.  What  color  changes  occur?  Dilute  some  of  the 
silver  sol  so  formed  with  water  and  describe  its  color.  What  reduces 
the  silver  oxide? 

Repeat  the  above  experiment,  using  a  drop  or  two  of  hydrazine 
hydrate  as  the  reducing  agent,  besides  gelatine. 

If  unsatisfactory  results  are  obtained,  repeat  the  experiment,  using 
smaller  amounts  of  AgNO3  solution  and  varying  other  conditions 
until  successful. 

Part  2.     Colloidal  Silver.     Method  of  Carey  Lea. 

Prepare  two  solutions  as  follows : 

Solution  A.    Mix:    10  per  cent  silver  nitrate  solution 20  cc. 

20  per  cent  Rochelle  salts  solution 20  cc. 

distilled  water 80  cc. 

Solution  B.    Mix:    30  per  cent  ferrous  sulphate  solution  .  .  10.7cc. 

20  per  cent  Rochelle  salts  solution 20  cc. 

distilled  water 80  cc. 

Pour  B  slowly  into  A,  stirring  rapidly.  The  solutions  must  be 
freshly  prepared  and  the  work  should  be  done  in  light  as  weak  as 
possible. 

Throw  out  the  precipitated  silver  by  means  of  a  centrifuge,  wash 
with  2  per  cent  Rochelle  salts  solution  and  again  separate  in  a  cen- 
trifuge. 

Obtain  a  camels-hair  brush  and  paint  some  of  the  silver  on  a  watch 
glass.  Dry  slowly  (without  heating  above  50°  C.)  and  note  the  color 
of  film  obtained. 

Place  a  crystal  of  iodine  in  the  center  of  the  yellow  silver  film. 
Record  all  that  happens.  Explain. 

References. 

Carey  Lea:  Am.  Jour.  Science,  (3)  37,  476  (1889);  38,  47,  129, 
237(1889);  41,179,259,482(1891);  Blake:  Zeit,  anorg.  chem., 
37,243(1903);  also  Svedberg:  Herstellung  (1909). 

Part  3.     Colloidal  Copper  (Gelatine  as  Protecting  Colloid.) 

Mix  equal  volumes  (5  cc.)  of  10  per  cent  gelatine  solution  (freshly 
prepared  and  warm)  and  5  per  cent  copper  acetate.  To  this  solution 

79 


add,  with  shaking,  a  very  slight  excess  of  sodium  hydroxide  (20  per 
cent).  A  purplish-blue,  clear  solution  should  result.  If  a  persistent 
precipitate  remains,  repeat  the  experiment,  using  a  more  concentrated 
gelatine  solution.  Perform  the  same  experiment,  using  5  cc.  of 
water  in  place  of  the  gelatine.  What  is  the  precipitate?  Does  it 
dissolve  in  an  excess  of  sodium  hydroxide? 

Heat  some  of  the  purplish-blue  copper  oxide-gelatine  solution  to 
boiling  and  add  a  few  drops  of  hydrazine  hydrate.  The  latter  is  a 
very  powerful  reducing  agent  and  will  reduce  the  oxide  to  metallic 
copper  in  alkaline  solution.  Continue  gently  to  heat  the  reaction 
mixture  until  a  dark,  blood-red  liquid  is  produced.  The  red  color 
is  due  to  finely  divided  copper.  Pour  some  of  the  liquid  into  water, 
noting  its  beautiful  color.  In  this  connection  cf.  Paal:  Ber.  35, 
2206,2219  (1902). 

EXPERIMENT  6 
Colloidal  Sulphur  (Condensation  Method) 

Reference.     Raff 6:     Kolloid-Zeit.,  2,  358  (1908);   8,  58  (1911). 

Place  a  cylinder  containing  70  grams  of  concentrated  sulphuric 
acid  (sp.  gr.  1.84)  in  ice  water  or  in  a  freezing  mixture  and  into  it  pour, 
drop  by  drop  and  with  constant  stirring  a  cold  solution  of  50  grams 
of  pure  crystallized  sodium  thiosulphate  in  30  cc.  of  distilled  water. 
Work  at  the  hoods,  as  H2S  and  SO2  are  given  off.  When  the  reaction 
is  complete,  transfer  the  mixture  to  a  beaker,  add  30  cc.  of  distilled 
water  and  warm  to  80°  on  a  water  bath  until  SO2  and  H2S  cease  to  be 
given  off.  Then  prepare  a  Buchner  funnel  and  filter,  connect  with 
the  suction  and  pour  in  hot  water  until  the  funnel  and  filter-flask  are 
warm.  Pour  out  this  wash  water  and  filter  the  hot  sulphur  hydrosol. 

Cool  the  warm  filtrate  in  ice  water  and  decant  the  supernatant  acid 
liquid.  To  some  of  the  precipitated  sulphur  add  water.  Is  it 
peptized? 

To  10  cc.  of  this  suspension  add  a  little  saturated  K2SO4.  What 
happens?  To  10  cc.  add  some  Na2SO4  solution.  Is  flocculation  so 
easy?  Note  difference  between  K2SO4  and  Na2SO4. 

Flocculate  some  of  the  sulphur  suspension  by  adding  a  soluble  salt 
of  potassium  and  allow  the  sulphur  to  settle.  Decant  the  super- 
natant liquid  and  wash  once  by  decantation.  Then  add  water  to 
the  precipitate  of  sulphur  and  shake  until  a  coarse  yellow  suspension 
of  sulphur  is  formed.  To  this  add  a  tiny  crystal  of  Na2SO4.  Con- 
tinue to  add  salt  cautiously  until  a  clear  yellow  suspension  of  sulphur 
is  formed.  What  is  this  process  called?  When  a  large  excess  of 
sodium  sulphate  is  added,  what  happens? 


SUB-GROUP  5 

EMULSIONS 

References.     Bancroft:   Jour.  Phys.  Chem.,  (1912-1918);  Briggs: 
Ibid.,  19,  210,  478  (1915);   24,  147  (1920). 

80 


EXPERIMENT  1 
Oil-in-Water  Emulsions 

Part  1.  In  a  150  cc.  glass  stoppered  bottle  place  45  cc.  of  benzene 
plus  5  cc.  of  1  per  cent  sodium  oleate  solution.  Then  shake  the 
bottle  and  contents  steadily  and  without  interruption  until  the  ben- 
zene is  completely  reduced  to  a  milk-white  emulsion  and  no  free 
benzene  remains  floating  at  the  surface.  Note  the  time  required  and 
the  approximate  number  of  shakes. 

Part  2.  Discard  the  emulsion  by  emptying  into  the  bottle  marked 
"benzene  residues"  and  repeat  the  experiment  making  a  change, 
however,  in  the  method  of  shaking.  Give  the  bottle  two  violent 
up  and  down  shakes  and  then  let  it  stand  on  the  desk  for  a  "rest 
interval"  of  about  thirty  seconds.  Continue  the  intermittent  shak- 
ing until  emulsion  is  completed.  Note  the  time  and  approximate 
number  of  shakes.  Compare  with  (1).  Explain. 

Part  3.  Again  discard  and  make  the  emulsion  in  still  another  way, 
as  follows : 

In  glass  stoppered  bottle,  place  2  cc.  of  sodium  oleate  solution  and 
to  this  add  1  cc.  of  benzene  from  a  burette.  Shake  thoroughly  until 
all  the  benzene  is  emulsified.  Then  add  another  cc.  of  benzene  and 
again  shake.  Repeat  this  process  until  about  100  cc.  of  benzene 
have  been  emulsified.  An  emulsion  having  the  consistency  and 
appearance  of  blanc-mange  should  result.  As  the  volume  of  emulsion 
increases,  more  benzene  may  be  added  each  time  before  shaking,  but 
if  too  much  is  added  the  emulsion  may  "break"  and  a  fresh  start 
become  necessary. 

Add  a  drop  of  HC1  to  some  of  this  emulsion.  What  happens? 
Explain. 

In  this  emulsion  the  oil  (benzene)  exists  in  drops  (disperse  phase) 
and  the  soap  solution  is  the  dispersion  medium. 

EXPERIMENT  2 
Water-in-Oil  Emulsions 

In  a  200  cc.  bottle,  as  in  the  previous  experiment,  place  10  cc.  of  a 
benzene  solution  of  magnesium  oleate.  Add  water  from  a  burette 
slowly  and  with  shaking,  following  a  procedure  similar  to  that  of  the 
preceding  experiment,  until  40  cc.  of  water  have  been  added.  How 
does  this  emulsion  compare  with  the  benzene-in-water  one?  In  this 
case  the  water  forms  the  drops  (disperse  phase)  and  the  soap  solution 
is  the  dispersion  medium.  This  may  be  proved  as  follows: 

Proof.  On  a  glass  plate  place  a  drop  of  water  and  with  a  glass  rod 
stir  in  some  of  the  emulsion  formed  in  Experiment  1 .  Does  it  mix 
easily?  On  another  portion  of  the  plate  place  a  drop  of  benzene  and 
stir  in  some  of  the  emulsion.  Does  it  mix  easily? 

Do  the  same  thing  with  some  of  the  emulsion  obtained  in  Ex'-eri- 
ment  2,  that  is,  stir  it  into  water  and  into  benzene. 

If  the  aqueous  liquid  is  the  outside  phase  the  emulsion  will  mix 
easily  with  water,  but  not  with  benzene.  The  reverse  is  true  when 
benzene  forms  the  outside  phase.  Newman:  Jour.  Phys.  Chem. 
13,35(1914). 

81 


EXPERIMENTAL  GROUP  XVII 

THERMOCHEMISTRY 

It  is  the  purpose  of  the  following  group  of  experiments  to  study  the 
thermal  effects  accompanying  chemical  action,  change  of  state  and 
similar  phenomena.  Many  instances  of  such  thermal  effects  have 
been  met  with  in  previous  experiments. 

References.  Thomsen  (Burke):  Thermochemistry  (1908). 
Thomsen:  Thermochemische  Untersuchungen  (1882-1886). 
Sackur  (Gibson) :  Thermochemistry  and  Thermodynamics  (1917) . 

Journal  articles. 

Mathews  and  Germann:  Jour.  Phys.  Chem.,  15,  73  (1911); 
Richards  and  Rowe:  Proc.  Amer.  Acad.,  43,  475  (1908) ;  Richards: 
Jour.  Am.  Chem.  Soc.,  31,  1275  (1909). 

Procedure  in  Laboratory.  F,  273-293  (1917);  OW,  119-138; 
T,  132-152. 

General  Directions. 

For  this  work  a  simple,  home-made  calorimeter  may  be  obtained 
from  the  Instructor. 

Two  special  thermometers  are  also  supplied.  These  must  be  com- 
pared with  each  other  in  the  usual  way  by  heating  in  a  well-stirred 
water-bath  between  10°  and  30°  C.  Number  each  thermometer  and 
reduce  all  subsequent  readings  of  temperature  to  readings  on  one  of 
these  thermometers. 

Having  assembled  the  calorimeter,  determine  the  water  equivalent 
by  experiment  several  times.  How  does  this  compare  with  the 
calculated  water  equivalent? 

Note.  Mix  weighed  and  approximately  equal  amounts  of  cold  and 
warm  water  so  that  the  final  temperature  of  the  mixture  is  about  equal 
to  that  of  the  room.  Weigh  out  water  to  grams  only  on  the  large 
balance. 

Report  the  water  equivalent  before  proceeding  with  the  experiments 
which  follow. 

EXPERIMENT  1 
Heat  of  Solution 

Part  1.  Qualitative.  Half  fill  a  test  tube  with  finely  powdered 
dry  NH4NO3  and  close  tube  with  a  rubber  stopper.  Then  add 
quickly  an  equal  volume  of  cold  water  and  mix  violently  to  produce 
instantaneous  solution.  Then  observe  the  temperature  of  the  solu- 
tion. Explain  the  extraordinary  drop  in  temperature.  How  does 

82 


this  method  of  making  a  freezing  mixture  compare  with  the  usual  one 
(ice-salt)  ?     Explain. 

Read  the  quaint  old  paper  on  this  subject  by  Robert  Boyle,  re- 
printed in  the  Philosophical  Transactions  of  the  Royal  Society  (Lon- 
don), 1,  86  (1666). 

Part  2.     Quantitative. 

Procedure.     T,  137. 

The  weighed  solute  is  introduced  into  a  known  amount  of  water 
contained  in  the  calorimeter.  A  convenient  method  is  to  make  a  thin 
walled  glass  bulb,  fill  it  with  the  solute  and  place  it  in  the  calorimeter. 
When  bulb  and  water  are  at  the  same  temperature,  break  the  glass 
and  allow  the  solute  to  dissolve  as  quickly  as  possible.  See  that  {he 
solute  is  very  finely  pulverized. 

Take  the  substance  assigned  from  the  following: 

(1)  NH4NO3  in  200  gram  molecules  of  water. 

(2)  KNO3  in  200  gram  molecules  of  water. 

(3)  NH4C1  in  200  gram  molecules  of  water. 

(4)  KC1  in  200  gram  molecules  of  water. 

Measure  the  heat  of  the  solution  and  derive  equation  (1)  below, 

Computations. 

S  =  p(a  +  w)     (ta  — ti)  (1) 

S  =  heat  of  solution  in  small  calories;  tt  =  initial  temperature  of 
water  and  bulb  in  calorimeter;  t2  =  final  temperature  when  solution 
is  complete;  a=  grams  of  water;,  w  =  water  equivalent;  1/p  = 
fraction  of  required  molecular  quantities  actually  used  experimen- 
tally. For  further  explanation  refer  to  Experiment  3  following. 


EXPERIMENT  2 

Heat  of  Dilution 
Procedure.     T,  139. 

In  this  experiment  the  solution  to  be  diluted  is  placed  in  the  upper 
vessel  and  the  water  is  placed  in  the  calorimeter.  The  solution  and 
water  are  then  mixed  and  the  thermal  affect  measured. 

Determine  the  heat  of  dilution  when  a  solution  represented  by 
NaCl  -f  10H2O  is  diluted  with  40  gram  molecules  of  water. 

Derive  equation  (2)  below. 

Computations. 

D  =  p  { (tf  —  tb)  [(a  +  b)  c  +  w]  —  (ta  —  tb)  (a  +  w) }        (2) 

D  =  heat  of  dilution  in  small  calories ;  ta  =  initial  temperature  of 
water;  tb  =  initial  temperature  of  solution;  tf  =  corrected  final 
temperature  of  mixture  whose  specific  heat  =  c ;  w  =  water  equiva- 
lent; a  =  grams  of  water;  b  =  grams  of  solution  to  be  diluted; 
1/p  =  fraction  of  required  molecular  quantities  actually  used 
experimentally. 

83 


EXPERIMENT  3 

Heat  of  Neutralization  of  'Acids  and  Bases 
Procedure.     T, 133. 

Place  the  acid  in  the  calorimeter  and"  the  base  in  the  upper  vessel. 
Mix  and  measure  the  heat  change. 

Computations. 

N  =  p  [b  (tf  —  tb)  +  (a  +  w)  (tf  —  ta)]  (3) 

N  =  heat  of  neutralization  in  small  calories;  ta  =  temperature  of 
acid;  tb  =  temperature  of  base ;  tf  =  temperature  of  mixture ;  a  = 
grams  of  water  contained  in  solution  of  acid;  b  =  grams  of  water 
contained  in  solution  of  base;  w  =  water  equivalent.  .  1/p  =  frac- 
tion of  required  molecular  quantities  used  experimentally.  Here  the 
specific  heat  of  the  mixture  is  assumed  to  be  unity. 
Derive  equation  (3). 

Part  1.     Sulphuric  Acid  and  Sodium  Hydroxide. 

Measure  the  heat  of  neutralization  for  each  of  the  following  cases: 

(a)  (2  NaOH  +  200  H2O)  +  (1/2  H2SO4  +  200  H2O). 

(b)  (2  NaOH  +  200  H2O)  +  (H2SO4  +  200  H2O). 

(c)  (2  NaOH  +  200  H2O)  +  (2  H2SO4  +  200  H2O). 

Part  2.     Phosphoric  Acid  and  Sodium  Hydroxide. 

(a)  (H3P04  +  200  H20)  +  (NaOH  +  200  H2O). 

(b)  (H3PO4  +  200  H2O)  +  (2  NaOH  +  200  H2O). 

(c)  (H3PO4  +  200  H20)  +  (6  NaOH  +  200  H2O). 

In  this  work  one  is  dealing  with  molecular  quantities  of  the  sub- 
stances involved.  For  instance  (2NaOH  +  200H2O)  means  80 
grams  of  NaOH  dissolved  in  3600  grams  of  H2O.  Again,  (1/2H2SO4 
+  200  grams  H2O)  means  49  grams  of  H2SO4  in  3600  grams  of  H2O. 
Obviously  such  volumes  of  acid  and  base  cannot  be  handled  con- 
veniently, so  one  chooses  some  convenient  fractional  part  of  the  acid 
and  base  solution,  for  example,  1/16  whence  1/p  =  1/16.  The  quan- 
tity of  the  solutions  to  use  in  the  case  of  H2SO4  and  NaOH  (Part  1) 
would  be  found  thus: 

1/16  (80  +  3600)  =  230  grams  of  the  NaOH  solution. 
1/16  (49  +  3600)  =  228  grams  of  the  H2SO4  solution. 
To  make  up  this  acid  solution  mix  3.06  grams  of  H2SO4  with .225 
grams  of  H2O. 

H2SO4  and  H3PO4  tables  may  be  found  in  the  Kalendar,  Vol.  1, 
and  elsewhere. 

EXPERIMENT  4 
Thermoneutrality  of  Salt  Solutions 

Measure  the  heat  change  when  solutions  of  the  following  pairs  are 
mixed: 

1 .  NaCl  +  200H2O  and  KNO3  +  200  H2O. 

2.  NH4C1  +  200H2O  and  KNO3  +  200  H2O. 

Take  the  pair  assigned,  placing   one  solution  in  the  upper  vessel 
and  the  other  in  the  calorimeter. 

84 


EXPERIMENTAL  GROUP  XVIII 

PHOTOCHEMISTRY 

The  purpose  of  the  following  experiments  is  to  study  qualitatively 
the  action  of  light  in  producing  and  accelerating  chemical  change. 

References. 

Bancroft:      Electrochemistry     of     Light,     Jour.     Phys.     Chem. 
(1908-1912). 

Bancroft:     Orig.  Comm.  8th  Int.  Cong.  App.  Chem.,  20, 31  (1912). 
Sheppard:     Photochemistry  (1914). 


EXPERIMENT  1 
Soluble  and  Insoluble  Sulphur 

Saturate  10  cc.  of  CS2  with  roll  sulphur.  Work  in  the  hood. 
Then  divide  into  three  portions  and  place  in  loosely  stoppered  test 
tubes. 

(a)  Expose  one  test  tube  to  direct  sunlight.     After  precipitation 
of  amorphous  sulphur  has  occurred,  set  aside  in  a  dark  place.     The 
amorphous  sulphur  will  dissolve.     It  may  be  necessary  to  wrap  the 
test  tube  in  dark  paper  to  protect  it  from  the  light. 

(b)  Place  another  portion  in  a  test  tube  which  is  immersed  in  a 
solution  of  CuSO4. 

(c)  Place  the  third  portion  in  a  test  tube  which  is  immersed  in  a 
solution  of  K2Cr2O7. 

Reference.     Rankin:     Jour.*Phys.  Chem  ,  11, 1  (1907). 

Note.  Do  not  stopper  the  test  tubes  too  tightly  when  exposing 
to  the  sunlight. 

Discussion. 

This  experiment  shows  in  a  very  satisfactory  way  how  light  dis- 
places the  equilibrium: 

Q  "  C 

soluble          insoluble. 

It  also  shows  that  the  activity  of  light  differs  for  different  wave 
lengths.  In  a  certain  sense  "light  is  a  mixture  of  reagents."  Light 
of  a  particular  wave  length  is  active  if  it  is  absorbed,  and  absorbed 
light  tends  to  shift  the  equilibrium  in  such  a  way  as  to  favor  the 
production  of  the  substance  which  absorbs  the  particular  light  less 
readily. 

85 


